Science: Honors Chemistry Flashcards
Summer work:
A quantity adopted as a standard of measurement
Unit
Summer work:
The temperature at which the minimum average kinetic energies of all particles occur
Absolute zero
Summer work:
The energy of an object that is due to the object’s motion
Kinetic energy
Summer work:
A change of matter from one form to another without a change in physical properties
Physical change
Summer work:
Any change in matter in which energy is released
Exothermic
Summer work:
A measure of the size of a body or region in 3-dimensional space
Volume
Summer work:
Elements that exist as single atoms
Monatomic
Summer work:
A measure of the average kinetic energy of the particles in an object, a measure of how hot or cold something is
Temperature
Summer work:
A substance or molecule that participates in a chemical reaction
Reactant
Summer work:
The energy transferred between objects that are at different temperatures. Energy is always transferred from higher-temperature objects to lower-temperature objects
Heat
Summer work:
A change that occurs when one or more substances change into entirely new substances with different properties
Chemical change
Summer work:
A substance that forms in a chemical reaction
Product
Summer work:
Changes in which the identity of a substance doesn’t change
Ex: state change
Physical changes
Summer work:
A measure of the amount of matter in an object. Not affected by the forces that act on the object
Mass
Summer work:
Consists of 2 or more atoms combined (bonded together) in a definite ratio, the smallest unit of a substance that keeps all of the physical and chemical properties of that substance
Molecule
Summer work:
The change of a substance from a liquid to a gas
Evaporation
Summer work:
The physical forms of matter: solid, liquid, gas, and plasma
States of matter
Summer work:
Law that states energy cannot be created or destroyed but can be changed from one form to another
Law of Conservation of Energy
Summer work:
The identities of substances change and new substances form
Chemical changes
Summer work:
A property of matter that describes a substance’s ability to participate in chemical reactions and can only be identified when an object tries to undergo a chemical change
Ex: flammability, reactivity w/ acid or oxygen
Chemical property
Summer work:
A characteristic of a substance that does not involve a chemical change
Ex: mass, color, texture
Physical property
Summer work:
A measure of the gravitational force exerted on an object (mass). Depends on gravity, varies with location.
Weight
Summer work: Evidence of chemical change: 1) evolution of a \_\_ 2) formation of a \_\_\_\_ 3) release or absorption of \_\_\_ 4) \_\_ \_\_\_ in the reaction system
Gas
Precipitate
Energy
Color change
Summer work:
Anything that has mass and takes up space
Matter
Summer work:
Any substance that has a definite composition
Chemical
Summer work:
The smallest unit of an element that maintains the properties of that element
Atom
Summer work:
A pure substance that are not elements made up of atoms of 2 or more different elements joined by chemical bonds
Compound
Summer work:
The quantity of heat required to raise a unit of mass of homogeneous material 1 K or 1*C in a specified way given constant pressure and volume
Specific heat
Summer work:
Something that has magnitude, size, or amount
Quantity
Summer work:
Any change in matter in which energy is absorbed
Endothermic
Summer work:
The process by which one or more substances change to produce one or more different substances
Chemical reactions
Summer work:
Pure substance that contain only one kind of atom that cannot be separated or broken down into simpler substances by chemical means
Element
Summer work:
_ L = ____ mL = ____ cm^3
1
1000
1000
Summer work:
Something that has a uniform structure or composition throughout, pure substances are distributed uniformly throughout the mixture
Homogeneous
Summer work:
A sample of matter that has definite chemical & physical properties
Pure substance
Summer work:
The capacity to do work. Cannot be created or destroyed.
Energy
Summer work:
Composed of dissimilar components, contains substances that are not evenly mixed
Heterogeneous
Summer work:
The mass of the object divided by its volume; often expressed as grams per cubic cm for solids and liquids and as grams per liter for gases
Density
Summer work:
Have fixed volume & shape
Have fixed volume but not a fixed shape
Have neither fixed volume nor fixed shape
Solids
Liquids
Gases
Summer work:
A sample of matter that contains 2 or more pure substances, a combination of 2 or more substances that are not chemically combined
Mixture
Stoichiometry:
Proportional relationship between 2 or more substances during a chemical reaction
Stoichiometry
Stoichiometry:
% yield= (____ yield/ ____ yield) x 100
Experimental
Theoretical
Stoichiometry:
The chemical which will effectively determine the amount of products that are formed & the chemical which will run out first as the reaction occurs and the reactants are consumed
Limiting Reagent
Stoichiometry:
The other reactants, the ones that are leftover when the limiting reagent runs out, are said to be _ __
In excess
Stoichiometry:
Coefficients in a balanced chemical equation show the relative # of moles of each substance in the reaction. You can use the coefficients as a conversion factor
Mole ratio
Gases:
The amount of force being exerted on a surface. Can be thought of as the weight of the atmosphere on top of an object.
Pressure
Gases:
1 atm = 14.7 psi
(___ __ __ ___)
Pounds per square inch
Gases:
The movement of individual molecules through a hole in a solid
Effusion
Gases:
Law of partial pressures equation
Total pressure= pressure #1 + pressure #2 (etc)
Gases:
For a fixed amount of an ideal gas kept at a fixed temperature, pressure and volume are inversely proportional
Boyle’s Law
Gases:
The atmosphere is __% Nitrogen, __% oxygen, __% Argon, <__% of other gases. 99.9% of Earth’s atmosphere can be found in the __ & ____
78% 21% 1% 1% Troposphere Stratosphere
Gases:
Random motion of small particles is called
Brownian motion
Gases:
The ratio between the pressure volume product and the temperature of a system remains constant
The combined gas law
Gases:
More collisions = __ pressure
Greater
Gases:
Gay-Lussac’s Law (pressure/temperature law) equation
P1/T1=P2/T2
Gases:
A 1 square inch column column of air weighs ~ __ lbs. This is defined as 1 ___ of pressure.
14.7
Atmosphere
Gases:
For gases collected over a liquid, the __ __ of the liquid must be accounted for. __ is present b/c some of the molecules are able to escape the surface of the liquid. The __ __ of the liquid is dependent on temperature
Partial pressure
Vapor
Vapor pressure
Gases:
Boyle noticed that pressure & volume of gases were __ __. As pressure went up, volume went __
Inversely proportional
Down
Gases:
Boyle’s Law equation
P1V1=P2V2
Gases:
Declared by John Dalton. States total pressure is the sum of the individual pressures added together
Law of partial pressures
Gases:
Melting & Boiling points are directly affected by: __ & __
Temperature
Pressure
Gases:
Molar mass equation
M=mRT/PV M: molar mass m: mass given R: .0821 T: temperature P: pressure V: volume
Gases:
Graham’s law equation
Rate 1/ Rate 2 = (molar mass 2/ molar mass 1)^.5
Gases:
The general movement of particles from an area of higher concentration to an area of lower concentration
Diffusion
Gases:
Density equation
D=MP/RT D: density M: molar mass P: pressure R: .0821 T: temperature
Gases:
Formula for pressure
P= F/A
P: pressure
F: force (lbs)
A: area (in^2)
Gases:
Einstein effectively proves the existence of atoms by explaining Brownian motion as the result of __ __ between the ___ & __ __
Random collisions
Particles
Gas molecules
Gases: Real gases are: 1) \_\_ point particles 2) \_\_ infinitely compressible 3) show at least a small amount of \_\_ to other molecules
Not
Not
Attraction
Gases:
For a fixed amount of an ideal gas kept at a fixed volume, pressure & temperature are directly proportional
Gay-Lussac’s Law
The pressure/temperature law
Gases:
Avogadro stated that equal volumes of gases contain __ number of particles
Equal
Gases:
Theory used to explain macroscopic properties of gases such as:
1)
2)
3)
Due to the __ of tiny particles (atoms/molecules)
Kinetic-Molecular Theory 1) pressure 2) temperature 3) volume Motion
Gases:
Ideal gas law equation
PV=nRT P: pressure (atm) V: volume (L) n: moles R: .0821 T: temperature (K)
Gases:
The combined gas law equation
(P1V1)/T1=(P2V2)/T2
Gases:
Gay-Lussac stated as the temperature of a gas increases, the volume __ __. Lussac called this __ __ b/c he used this person’s stuff
Increases proportionally
Charles’s law
Gases:
For a fixed amount of an ideal gas kept at a fixed pressure, volume and temperature are directly proportional
Charles’s Law
Gases:
Gas pressure can be measured using a __
Manometer
Gases: Properties of gases 1) great \_\_ b/w particles 2) is a \_\_\_ 3) low \_\_\_ 4) highly \_\_\_\_ 5) \_\_ fill a container
distance fluid density compressible Completely
Gases: Ideal gases consist of: 1) random moving \_\_ \_\_ that 2) are \_\_ compressible and 3) have no intermolecular \_\_\_
Point particles
Infinitely
Attractions
Gases:
Charles’s law equation
V1/T1=V2/T2
Gases:
Avogadro’s Law equation
V1/n1=V2/n2
Gases:
Gay-Lussac stated as the temperature of a gas increases, the pressure __ __
Increases proportionally
Gases:
For an ideal gas kept at a fixed pressure and temperature, volume and moles are directly proportional
Avogadro’s Law
Solutions:
Solutions must be a __ mixture
Homogeneous
Solutions:
A solution that cannot dissolve any more solute under the given conditions
Saturated solution
Solutions:
When ionic compounds dissolve in water, they ___. Compounds that break up into ions when dissolved are called ____
Dissociate
Electrolytes
Solutions:
Calculating concentration equation
% conc = ( solute / total solution ) * 100
Solutions:
Solutions can be in the phases of __, ___, & ___
Solids
Liquids
Gases
Solutions:
Solubilities of gases generally __ with increasing ___, but ___ with increasing ___
Decrease
Temperature
Increase
Pressure
Solutions:
Polar molecules dissolve __ molecules. Non polar molecules dissolve __ molecules
Polar
Non polar
Solutions:
A solution holding more dissolved solute than what is required to reach equilibrium at a given temperature
Supersaturated solution
Solutions:
The concentration of a solution expressed in moles of solute / liters of solution. Abbreviated __
Molarity
M
Solutions:
A typical solution consists of a __ solute dissolved in a __ solvent. When both substances are in the same ___, the substance that there is the most of is the ____.
Solid
Liquid
Phase
Solvent
Solutions:
When a substance is written with square brackets, [ ], = “__ __ __”. M pronounced “__” when used in a sentence
The concentration of
Molar
Solutions:
Ability of one substance to dissolve into another at a given ___ & ___. Expressed in terms of the amount of __ that will dissolve in a given amount of __ to produce a __ __
Solubility Temperature & pressure Solute Solvent Saturated solution
Solutions:
Dilution equation
M1V1 = M2V2
M: molarity
V: volume
Solutions:
Solubilities of solids generally __ with increasing __, but generally ___ with increasing ___
Increase
Temperature
Remain constant
Pressure
Solutions:
Solids that dissolve but don’t dissociate. Non-ionic solids that will dissolve in water, but will not conduct electricity
Non-electrolytes
Solutions:
Dissolved salts. Ionic solids that will dissolve in water and conduct electricity
Electrolytes
Solutions:
A solution that contains less solute than a saturated solution and that is able to dissolve additional solute
Unsaturated solution
Solutions:
Saturated solutions occur when __ is reached between __ & __ of the __. This is know as __ __
Equilibrium Dissolving Recrystallizing Solute Dynamic Equilibrium
Solutions:
ppm= __ __ __ or __ __ / __ __
ppb= __ __ __ or __ __ / __ __
Parts per million
mg solute / kg solution
Parts per billion
ug solute / kg solution
Solutions:
Solutions have two parts:
__: substance being dissolved (changes phases)
__: substance doing the dissolving
Solute
Solvent
Equilibrium:
Reactions where almost all of the reactants are turned into products. Very little, if any reactant remains when the reaction is complete
Completion reactions
Equilibrium:
State of balance in which the rate of a forward reaction equals the rate of the reverse reactions and the concentrations of products and reactants remain unchanged. It is a state of __ ___
Chemical equilibrium
Dynamic equilibrium
Equilibrium:
Effect of change on equilibrium (describe the shift):
Exothermic:
Lowering temperature
Shift right
Equilibrium:
If Keq >1, then the reaction is considered __, because there will be more __ than __ at equilibrium
Favorable
Product
Reactants
Equilibrium:
Effect of change on equilibrium (describe the shift):
Concentration:
Add
Shift to opposite side
Equilibrium:
Effect of change on equilibrium (describe the shift):
Exothermic:
Raising temperature
Shift left
Equilibrium: The system responds to 3 different kinds of stress: 1) changes in 2) changes in 3) changes in
Concentrations of reactants or products
Temperature
Pressure
Equilibrium:
Effect of change on equilibrium (describe the shift):
Increase in pressure
Shift to side with fewer gas molecules
Equilibrium:
States that a system in equilibrium will oppose a change in a way that helps eliminate the change
Le Chatelier’s Principle
Equilibrium:
Effect of change on equilibrium (describe the shift):
Endothermic:
Raising temperature
Shift right
Equilibrium:
For Keq constant, do not include ___ or __ as their concentrations don’t change
Solids
Liquids
Equilibrium:
If Keq=1, then neither reactants nor products are ___, both exist in __ __ at equilibrium
Favored
Equal amounts
Equilibrium:
Reduction of the solubility of a salt in the solution due to the addition of a common ion
Common ion effect
Equilibrium:
The Equilibrium constant equation
Keq= [products]^x / [reactants]^y
Where [ ] is the concentration in molarity of each substance and its coefficient is the exponent
Equilibrium:
Is Keq for the special case of a slightly soluble salt in water
Ksp
Equilibrium:
Effect of change on equilibrium (describe the shift):
Endothermic:
Lowering temperature
Shift left
Equilibrium:
Effect of change on equilibrium (describe the shift):
Decrease pressure
Shift to side with more gas molecules
Equilibrium:
A reaction in which the products re-form the original reactants
Reversible reactions
Equilibrium:
If Keq <1, then the reaction is considered __, because there will be more ___ than ___ at equilibrium
Unfavorable
Reactants
Products
Equilibrium:
Change H < 0 = __ reaction
Change H > 0 = __ reaction
Exothermic
Endothermic
Equilibrium:
Effect of change on equilibrium (describe the shift):
Concentration:
Subtract
Shift to same side
Acids & Bases:
Calculating pOH & [OH-]
pOH = -log[OH-] [OH-]= 10^-pOH
Acids & Bases:
__ have a pH < 7, __ have a pH > 7
Acids
Bases
Acids & Bases:
Only a small fraction of its molecules are ionized at any given time, producing only a few hydronium ions
Weak acids
Acids & Bases:
Arrhenius definition of a base
Any substance that forms hydroxide ions in solution
Acids & Bases:
Compounds that can reversibly change color depending on the pH of the solution or other chemical change
Indicators
Acids & Bases:
A solution is considered __ if the concentrations of the hydronium & hydroxide ions are equal (pH of 7)
Neutral
Acids & Bases:
Value used to express the acidity or basicity of a solution. It is defined as the logarithm of the reciprocal of the concentration of hydronium ions
pH
Acids & Bases:
Arrhenius definition of an acid
Any substance that forms hydronium ions in solution
Acids & Bases:
Self-ionization constant of water
Kw
Acids & Bases:
Releases few hydroxide ions in solution
Weak bases
Acids & Bases: Properties of acids: 1) Solutions \_\_ \_\_ well 2) react with many \_\_ 3) \_\_ taste 4) acids generate \_\_ \_\_ (\_\_) in water
Conduct electricity
Metals
Sour
Hydronium ions (H3O+)
Acids & Bases:
What does pH stand for?
Power of hydrogen
Acids & Bases: Self Ionization of water Equilibrium equation: [H3O+]=[OH-]= \_\_\_\_ \_\_\_\_ Kw = [\_\_][\_\_]
H2O (base) + H2O (acid) <-> H3O+ + OH-
1x10^-7 M
1.00x10-14=[H3O+][OH-]
Acids & Bases:
pH equation
pH = -log [H3O+] or pH = -log[H+] [H3O+]= 10^-pH or [H+]= 10^-pH
Acids & Bases:
Examples of indicators
Phenolphthalein, thymol blue, pH paper, universal indicator
Acids & Bases:
Describes a substance, such as water, that has the properties of an acid and a base
Amphoteric
Acids & Bases:
Ionizes completely in a solvent, producing a lot of hydroxide ions
Strong bases
Acids & Bases:
pH+pOH=__
14
Acids & Bases:
The __ of the two ion concentrations is always equal to __
Product
Kw
Acids & Bases:
A substance that donates a proton to another substance
Bronsted-Lowry Acids
Acids & Bases:
[H+]> 1x10^-7: ___
[H+]< 1x10^-7: ___
Acid
Base
Acids & Bases:
Ionize completely in water, producing a lot of hydronium ions
Strong acids
Acids & Bases:
A substance that accepts a proton from another substance
Bronsted-Lowry Bases
Acids & Bases: Properties of Bases: 1) \_\_ 2) \_\_ taste 3) \_\_ to the touch 4) generate \_\_ \_\_ (\_\_) in water
Electrolytes
Bitter
Slippery
Hydroxide ions (OH-)
Atomic Theory:
English chemist / physicist / meteorologist:
Many areas of influence:
The behavior of gases, interactions of light & heat, and the behavior of chemicals
John Dalton
Atomic Theory:
Discovered negatively charged particles (electrons). Proposed the “plum pudding” model. Postulated positive particles. Did the cathode ray experiment
JJ Thomson
Atomic Theory:
Aimed alpha radiation at light elements like boron. Found it gave off an extremely penetrating radiation. Thought it produced high energy gamma rays.
Walther Bothe & Herbert Becker
Atomic Theory:
All atoms of a given element must have the same number of ____. In order to be neutral, atoms must have the same number of ___ as ___ (charges must cancel out to 0)
Protons
Electrons
Protons
Atomic Theory:
The sum of the protons and neutrons in an atom equals the ___
Mass number
Atomic Theory:
Updated atomic model:
Discarded __ __ model
Electrons must orbit around central ____ -> Planetary model
Nearly all of the mass is located in the dense, central, positively charged ____
Plum pudding
Nucleus
Nucleus
Atomic Theory:
Positively charged, easily stopped
Negatively charged, stopped by a sheet of aluminum
High energy light (no charge), penetrates a lot of material
Alpha radiation
Beta radiation
Gamma radiation
Atomic Theory:
Said matter is composed of indivisible parts. The parts are called atomos. Properties of matter are due to the size, shape, and weight of the atomos. His idea was not popular and was abandoned for over 2000 years.
Democritus
Atomic Theory:
Average mass of all naturally occurring isotopes. Provides more accurate mass of a typical sample. A weighted average.
Atomic mass (weight)
Atomic Theory:
____ of an atom is sum of its protons, neutrons, & electrons. Measured in ____.
Mass
Daltons
Atomic Theory:
Matter can neither be created or destroyed (in a chemical reaction), but only changed from one form to another. Total mass of reactants must equal total mass of products.
Law of Conservation of Mass
Atomic Theory:
Experiment that discovered the electron. Gas filled tube with metal plates with a positive end (anode) and a negative end (cathode).
Cathode ray experiment
Atomic Theory:
Said you can continue to split a substance in half until you broke it down into the elements. (Elements could be divided in half indefinitely)
Aristotle
Atomic Theory:
Continued work of Bothe & Becker: Aimed Bothe’s new beam at paraffin, ejected high energy protons, misinterpreted results.
Died from radiation exposure.
Irene Joliot-Curie
Atomic Theory:
When two or more different compounds are made up of the same elements, their percent compositions will always be different.
Law of Multiple Proportions
Atomic Theory:
Shot a beam of particles to a piece of gold foil. Most particles pass straight through foil but some particles are scattered. Showed plum pudding model was incorrect. Created the Rutherford model.
Gold foil experiment
Atomic Theory:
Did not believe the work of Joliot-Curie (she said beam was light waves). Discovered neutron (no charge, mass slightly more than a proton)
James Chadwick
Atomic Theory:
Discovered the positively charged dense central portion of the atom (nucleus) using his “gold foil experiment”. Atoms are mostly empty space. Discovered the proton.
Ernest Rutherford
Atomic Theory:
Danish physicist. Worked on the electron/nucleus problem. Father of Quantum Mechanics. Solution: electrons are located in specific ___ __. Electrons move in a ___ orbit around the nucleus. Areas between orbits are __ ___
Niels Bohr
Energy levels
Definite
Not allowed
Atomic Theory:
Heavy, positive charge, repelled by the nucleus
Nearly massless, negative charge, repelled by the electrons surrounding an atom
Heavy, no charge, no interactions with nucleus or electrons (can pass through a lot of material, if fast enough can break nucleus apart)
Protons
Electrons
Neutrons
Atomic Theory:
A compound always contains the same elements in exactly the same proportions by weight. Compounds will always have the exact same percent composition. It is intrinsic (doesn’t matter how much you have, the percentage will always be the same)
Law of Definite Proportions
Atomic Theory:
Atoms can have more or less ___. ____: atoms with the same # of protons, but different # of neutrons. Isotopes occur in different proportions in nature.
Neutrons
Isotopes
Atomic Theory: Dalton's Atomic Theory: 1) Matter is composed of \_\_\_ 2) All atoms of a given element are \_\_\_ 3) Properties of different elements are \_\_\_\_\_ 4) Atoms combine in \_\_\_\_ \_\_\_\_ 5) We cannot \_\_\_ or \_\_\_ atoms
Atoms Identical Different Whole numbers Create, destroy
Atomic Theory:
Explanation that is testable by doing experiments
Observation of something with rules that we follow without explanations
Theory
Law
Periodic Trends:
Describe the halogens
Group 17 [np^5] All nonmetals Gases (F,Cl); liquid (Br); solids (I,At) Name means "salt former" Very toxic! React with sodium to form a salt with a 1-to-1 ratio
Periodic Trends:
Rearranged table according to electronic charge in 1914. Became # of protons after 1918. Noticed his new table had spots for #’s 43, 61, 72 & 75. Produced the modern periodic table we know today.
Henry Moseley
Periodic Trends:
Describe the lanthanoids
1st row on bottom of table [4f^x]
AKA Lanthanides & rare earths
Ce is not so rare (25th most abundant)
So similar that they are very difficult to separate (Moseley)
Most deflect UV- used in sunglasses
Shiny, silvery white, soft, react violently with most nonmetals, tarnish in air
Periodic Trends:
The higher the electronegativity, the ____ the atoms ____ for electrons
Harder
Pull
Periodic Trends:
What two elements are usually found as 8 atoms?
What element is usually found as 4 atoms?
Sulfur and Selenium
Phosphorus
Periodic Trends:
Measure of how badly an element wants to gain an electron
Electronegativity
Periodic Trends:
What element has the highest electronegativity?
What element has the lowest electronegativity?
Fluorine
Francium
Periodic Trends:
A horizontal row of elements in the periodic table. Have the same # of occupied ___ ___. ___ similar at all.
Period
Energy levels
Not
Periodic Trends:
Electronegativity ____ from left to right
Electronegativity ____ from top to bottom
Increases
Decreases
Periodic Trends:
The amount of energy required to remove an electron
Ionization energy
Periodic Trends:
What are the 4 groups of representative elements?
Alkali Metals (Group 1)
Alkaline Earth Metals (Group 2)
Halogens (Group 17)
Noble Gases (Group 18)
Periodic Trends:
Describe the noble gases
Group 18 [np^6]
Un reactive gases- colorless & odorless
Some of the last natural elements to be discovered
Once called “inert gases”
Monatomic in nature (1 atom)
Don’t combine with other atoms because their outer p-orbital are full
Periodic Trends:
When 2 atoms have the same number of electrons but a different amount of protons, the one with ___ protons will pull the electrons in tighter and make the atom ____
More
Smaller
Periodic Trends:
What are some common salts of alkali metals?
LiCl NaCl KCl RbCl CsCl FrCl
Periodic Trends:
A vertical column is called a ___ or ____. Have ____ properties
Group
Family
Similar
Periodic Trends:
What are the 4 trends in the periodic table?
Atomic Radius
Electronegativity
Ionization Energy
Reactivity
Periodic Trends:
The atomic radius ______ from left to right
The atomic radius _____ from top to bottom
Decreases
Increases
Periodic Trends:
Most reactive metals are ____ ___ and to the ____
Most reactive nonmetals are ____ ___ and to the ____
Further down
Left
Higher up
Right
Periodic Trends:
What element has the highest ionization energy?
What element has the lowest ionization energy?
Fluorine
Francium
Periodic Trends:
Atoms in their _____ ____ tend to be more dangerous/poisonous than those in their ____ ____
(Exceptions: __ & __)
Elemental state
Reacted state
Cu & Pb
Periodic Trends:
Typically placed in its own category. Most of the time, behaves like a nonmetal. Under certain circumstances, can behave like a metal.
Hydrogen
Periodic Trends:
What are the groups that are non-representatives?
Groups 13-16 (Boron, Carbon, Nitrogen, Oxygen): have more differences than similarities
Transition metals
Lanthanoids
Actinoids
Periodic Trends:
What are properties of metals?
Lustrous (shiny); malleable (pounded into thin sheets); ductile (pulled into wires); conductive (heat & electricity); form solid oxides when burned; tend to react with acids to form hydrogen gas
Periodic Trends:
What are some common salts of alkali earth metals?
BeCl,2 MgCl,2 CaCl,2 SrCl,2 BaCl,2 RaCl,2
Periodic Trends:
Exhibit nearly perfect periodicity. All members of these groups behave as expected.
Representative elements
Periodic Trends:
What were some problems with Mendeleev’s table?
Based on atomic mass, had to switch a few elements (tellurium & iodine) to keep reactivities in order.
Believed he predicted too many elements.
Used it for half a century
Periodic Trends:
Describe the alkaline earth metals
Group 2 [ns^2]
Harder & denser then alkali metals
Lustrous, oxidize slowly when exposed to air
React with water or steam to form a base
React with chlorine to form a salt with 1-to-2 ratio
Periodic Trends:
How is the modern periodic table set up?
A chart of the elements showing the repeating pattern of their properties. Elements are arranged in rows and columns by increasing atomic #. Atomic # increases by 1 between each element.
Periodic Trends:
Arranged known elements in a table by atomic mass in 1863. Noticed a repeating pattern every 8th element in 1865. Was laughed at by peers.
John Newlands
Periodic Trends:
What are some common salts of halogens?
NaF NaCl NaBr NaI NaAt
Periodic Trends:
Describe the transition metals
Groups 3 to 12 [nd^x] Central portion of periodic table Behavior & appearance vary Have a variable oxidation state (charge) Different oxidation states that can produce different colors Often used to make pigments
Periodic Trends:
_______ is by far most abundant
(4 out of every 5 atoms in the universe)
Hydrogen
Periodic Trends:
What are properties of nonmetals?
Tend to be:
Dull; brittle (when solid); insulators; form gaseous oxides; don’t react much with acids; have lower melting & boiling points (than metals)
Periodic Trends:
Describe metalloids
AKA: “semi-metals” or “staircase elements”
Combination of properties of metals & nonmetals.
Boron, silicon, germanium, arsenic, antimony, tellurium, & polonium
Periodic Trends:
What did Moseley’s new periodic table do?
Gave experimental meaning to atomic #. Gave reason for tellurium & iodine being switched. Easily separated rare earth metals. Used to predict how many elements remained between others.
Periodic Trends:
7 elements always form _____ molecules when they are isolated in their elemental state (ALWAYS!).
Means two atoms
Two atoms of the same element bond together.
5 gases, 1 liquid, & 1 solid.
Diatomics
Periodic Trends:
Why does the atomic radius increase from top to bottom? Why does the atomic radius decrease from left to right?
More energy levels to fill
More charge, attraction pulls them in tighter
Periodic Trends:
What are the names of all groups/families?
Alkali metals Alkaline earth metals Halogens Nobles gases Transition metals Inner transition metals- lanthanoids (rare earths), Actinoids
Periodic Trends:
Atoms are less dangerous as ___
Ions
Periodic Trends:
Produced a more orderly table independent of Newlands’ work in 1869. Also used atomic mass. Left blanks for yet-undiscovered elements. Predicted properties of Ga, Sc, and Ge which were discovered in 1875,77, & 86. Credited with developing the periodic table.
Dmitri Mendeleev
Periodic Trends:
Describe the Alkali metals
Group 1 [ns^1]
Soft, lustrous, oxidize when exposed to air
Difficult to isolate- never found in nature
React (violently) with water to form a base
React with chlorine to form a salt with a 1-to-1 ratio
Periodic Trends:
Other families have similarities, but do not behave exactly as expected. Some lumped together for other reasons.
Non-representatives
Periodic Trends:
Chemical properties repeat every 8 elements
Law of Octaves
Periodic Trends:
When an atom changes to a(n):
Anion- it gets _____
Cation- it gets _____
Bigger
Smaller
Periodic Trends:
Which element has the largest atomic radius?
Which element has the smallest atomic radius?
Francium
Hydrogen or Helium
Periodic Trends:
Ionization energy _____ from left to right
Ionization energy _____ from top to bottom
Increases
Decreases
Periodic Trends:
Elements are found on different parts of the periodic table:
____ to the left (majority of elements)
____ to the right (18 elements)
____ found on a “staircase” dividing metals & nonmetals (7 elements)
____&____(metals) added to bottom to make table manageable
Metals
Nonmetals
Metalloids
Lanthanoids & Actinoids
Periodic Trends:
The reduction of the attractive force between a positively charged nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charges of the inner electrons
Electron shielding
Periodic Trends:
Describe the Actinoids
2nd row on bottom of table [5f^x] AKA Actinides All are radioactive Not as similar as the lanthanoids Only Th and U are common in nature Most are man-made Nuclear fallout Particle colliders
Periodic Trends:
What are the 7 diatomic elements?
Hydrogen Oxygen Fluorine Bromine Iodine Nitrogen Chlorine
Quantum Mechanics: s orbitals: \_\_\_\_, larger n = bigger \_\_\_ p orbitals: \_\_\_ \_\_\_, larger n = bigger \_\_\_\_ d orbitals: only possible when n>_ f orbitals: only possible when n>_
Spherical Sphere Dumbbell shape Dumbbell 2 3
Quantum Mechanics:
4 quantum numbers (specify location of electrons in atom):
1) n (principle) = ___ ___ (n>0). Lower the #, closer to the nucleus
2) l (azimuthal) = ___ ___. __ are abbreviated s, p, d, f
3) m = _____ ____. s=1, p=3, d=5, f=7
4) m(s) = ____. Opposite directions
Energy level
Orbital shape, Shapes
Suborbital orientation
Spin
Quantum Mechanics:
The Dreaded Double Slit:
When light passes through 2 narrow slits, it creates interference patterns just like ___ on a pond. We see the same interference patterns when we shoot ___ at a double slit. Electrons are ___ & __.
Waves
Electrons
Particles
Waves
Quantum Mechanics:
Add one electron in each suborbital before pairing up
Hund’s Rule
Quantum Mechanics:
A measure of how often a wave passes a point. Waves per second = ____ (Hz)
Means “per second”
Frequency
Hertz
Quantum Mechanics:
J: ___, unit for energy
Joules
Quantum Mechanics:
Performed the double slit experiment & proved light existed as a wave.
Thomas Young
Quantum Mechanics:
Particle of light
Photon
Quantum Mechanics:
Research with black-body radiation led to the idea that energy is quantized. Energy of an oscillator is strictly proportional to its frequency. As frequency increases, energy of photon (quanta) increases and vise versa.
Max Planck
Quantum Mechanics:
___ is constant (c)
Speed of light
Quantum Mechanics:
The more you know about the position of a particle, the less you know about its momentum (speed and direction) and vice versa. The act of measuring causes an inherent ____ in the universe!
Heisenberg’s Uncertainty Principle
Uncertainty
Quantum Mechanics:
Electrons fill orbitals with the ____ energy first. s can have _ electrons. p can have _ electrons. d can have _ electrons. f can have _ electrons. Electron orbital diagrams created by drawing a box for each orbital. Put electrons in one-by-one, obeying ____ ___
Lowest 2 6 10 14 Hund's Rule
Quantum Mechanics:
No 2 electrons can have the same 4 quantum numbers. Electrons cannot stack up on each other. If an orbital is full, the next electron must go to a higher orbital. This is what makes matter solid. Each electron must have a different set of quantum #’s.
Pauli Exclusion Principle
Quantum Mechanics:
Louis de Broglie tells us why we have specific orbits. Only standing waves can exist within a ___
Space
Quantum Mechanics:
Shorter wavelengths have ___ energies and frequency
Longer wavelengths have ___ energies and frequency
Higher
Lower
Quantum Mechanics:
Take a cat. Put it in a box with a device that may or may not kill it. Is the cat dead or alive? It is both! The cat exists in a “superposition” (a state between dead and alive) until you check, then the wave function collapses and the cat is dead or alive (but not both). The idea has been proven.
Schrodinger’s Cat
Quantum Mechanics:
Neils Bohr proposed:
Electrons are limited to ___ ___ around the nucleus
Electrons become ___ (absorb energy) and “jump” to a ___ ____
Specific orbits
Excited
Higher orbital
Quantum Mechanics: \_\_\_ are in orbitals. Orbitals differ in: 1) 2) 3) Electrons have \_\_\_
Electrons Size Shape Orientation Spins
Quantum Mechanics:
Described light via the corpuscular theory: Light is composed of tiny, discreet packets called “corpuscles”. More influential, so his theory won out.
Sir Isaac Newton
Quantum Mechanics:
Since different wavelengths are emitted, we can calculate the difference in ___ between ___
Energy
Orbitals
Quantum Mechanics:
Described light using a wave theory. Light propagates through space just like ripples on a pond.
Christiaan Hyugens
Quantum Mechanics:
___ & ___ are inversely related. When one increases, the other must decrease. Waves with __ wavelengths have __ frequencies.
Waves with __ wavelengths have __ frequencies
Wavelength & frequency
Long, low
Short, high
Quantum Mechanics:
Heisenberg killed _____ & Schrodinger’s wave equation brought in probabilities. There are no absolute answers, only ___
Determinism
Probabilities
Quantum Mechanics:
Albert Einstein proved light acts as a particle (photon). Light was already proven as a wave. Light behaves like a wave and a particle at the same time. Depends on how/what you are measuring
Wave-Particle Duality
Quantum Mechanics:
Height of a wave:
Highest point of a wave:
Lowest point of a wave:
Amplitude
Peak
Trough
Quantum Mechanics:
Distance between identical sections on a wave. Measured in meters
Wavelength
Quantum Mechanics:
Photons of different wavelengths have different ___. Calculate energy using E= hv
h= ____ ____ : 6.626 x 10^-34 J•s
Energies
Planck’s constant
Quantum Mechanics: s orbitals: _ possibility p orbitals: _ possibilities d orbitals: _ possibilities f orbitals: _ possibilities
1
3
5
7
Quantum Mechanics:
Atoms are “built up”. Add a proton to the nucleus. Add an electron to the lowest possible orbital.
Aufbau Principle
Quantum Mechanics:
Unsatisfied with Heisenberg’s mathematical picture, came up with a wave-like model. All matter can be described as a wave function. All positions are equally possible, but not all positions are equally probable
Erwin Schrodinger
Quantum Mechanics:
All electromagnetic waves travel at the same ___ in a vacuum (space).
Speed of light = c = ___ x 10* m/s
Usually rounded to __ x 10* m/s
Wave speed
Speed
2.998
3.0
Quantum Mechanics:
When gaseous atoms/molecules are heated, they emit ___. __ is specific to each substance. Spectral analysis shows discrete ___
Light
Color
Lines
Quantum Mechanics:
According to his wave equation, all possible outcomes exist at one time until you ___ it! The act of ____ causes one (and only one) result to “pop” into existence.
Measure
Measuring
Quantum Mechanics:
Bohr’s Model:
Electrons only permitted in ___ orbits. The __ state is the open orbit closest to nucleus (lowest available energy). The __ state is an orbit farther away from nucleus (it has higher energy). ___ __ # (n) is given to determine orbit. Small n means a small ___ (closer to nucleus). n>0
Circular Ground Excited Principal Quantum Radius
Quantum Mechanics:
A rhythmic disturbance that carries energy through matter or space. Travels at a constant speed through a medium. Can change speed if the medium changes. Speed is product of ____ & _____
Waves
Wavelengths
Frequency
Quantum Mechanics:
Said if light can behave as both a wave and a particle, then it reasons that things like electrons can as well
Louis de Broglie
Quantum Mechanics:
One consequence of the uncertainty principle is that the electron could no longer be considered as in an ___ location in its orbital. The electron has to be described by every point where the electron could possibly ____
Exact
Inhabit
Quantum Mechanics:
Occurs when electrons absorb energy and jump to a higher energy level and when electrons lose energy and fall back down to a lower orbital. When electrons lose energy, the emit a photon of light.
Light emission
Quantum Mechanics: 3 properties of waves: 1) \_\_\_ \_\_\_: when 2 waves come in contact with each other. Can be either: a) \_\_\_\_: waves amplitudes combine and doubles. b) \_\_\_\_: waves cancel out 2) \_\_\_\_: spreading out beyond a barrier 3) \_\_\_\_: bending of light
Wave interference Constructive Destructive Defraction Refraction
Bonding:
___ form when atoms attempt to fill their outer s & p orbitals to be ______ with the noble gases
Ions
Isoelectronic
Bonding:
Atoms with ____ electronegativity will share electrons
Similar
Bonding:
Ions are ___ ___ & have the same configuration as a ___ ___, but ions have a ____ and ____ ____ when dissolved in ____, unlike the ____ ____
More stable Noble gas Charge Conduct electricity Water Noble gases
Bonding:
Form from metals
Can be ___ metals or ___
Metallic substance
Pure
Alloy
Bonding:
Aluminum tends to form ions with a ____ charge
Zinc tends to form ions with a ____ charge
Silver tends to form ions with a ____ charge
+3
+2
+1
Bonding:
Prefixes 1-10
1) mono
2) di
3) tri
4) tetra
5) penta
6) hexa
7) hepta
8) octo
9) nona
10) deca
Bonding:
Bonds form between atoms due to the interactions of _____
Electrons
Bonding:
Properties of ions are _____ ____ from their neutral counterpart
Completely different
Bonding:
Form between a nonmetal and another nonmetal
Share electrons
Covalent compound
Bonding:
Electrons closest to the nucleus
Unable to participate in bonding
Shielded by electrons farther away
Core electrons
Bonding:
Properties are dictated by the ___ & ___ of an element’s ___
Number
Configuration
Electrons
Bonding:
Atoms will attempt to gain or lose electrons to have 8 electrons in their outer s & p orbitals
Octet Rule
Bonding:
Ions made up of more than 1 atom
In ______ ions, ____ ___ form between atoms within the ion
Electrons have been ___ or ____ during formation
Polyatomic Polyatomic Covalent bonds Gained Lost
Bonding: Metallic substances: Usually \_\_\_, but a few \_\_\_\_ at or near room temperature Will not dissolve in \_\_\_ or \_\_\_ All metals \_\_\_ \_\_\_ at least a little bit
Solid, liquids
Water, oil
Conduct electricity
Bonding:
Roman numerals 1-10
I II III IV V VI VII VIII IX X
Bonding:
Few bonds are strictly ___ or ____
Ionic
Covalent
Bonding:
All compounds are neutral, so the total charge must add up to _
0
Bonding:
& block elements almost always obey the octet rule
& block elements do not always follow the octet rule
s p
d f
Bonding: All ionic compounds are called \_\_\_ They have \_\_\_ \_\_\_ melting points All are \_\_\_ at room temperature Will \_\_\_ in water (if they \_\_\_\_)
Salts Very high Solid Dissociate Dissolve
Bonding:
When writing formulas containing polyatomic ions, it is important to remember that the _____ are part of the polyatomic ion
Subscripts
Bonding:
Ions made up of single atoms
Ex: Mg^+2, Br^-1
Monatomic ions
Bonding:
When the outer s & p orbitals are full, the atom is very _____
Stable
Bonding:
Electrons farthest from the nucleus
Determine the chemical reactivity of the atom
Valence electrons
Bonding:
This means you must have an equal # of _ & _ charges
+
-
Bonding:
Anions are named by dropping the ___ ____ of the element and adding “-__”
Chlorine becomes ___
Oxygen becomes ____
Final syllable
-ide
Chloride
Oxide
Bonding:
Change ending of the second element to “-ide”
Use prefixes to indicate the # of atoms each element present
Rules for naming covalent compounds
Bonding:
The unequal sharing of electrons within a bond leads to the formation of an ___ ____
Electric dipole
Bonding:
Cations form when atoms ____ electrons, leaving more ____ than ____
All ___ form cations
Lose
Protons
Electrons
Metals
Bonding: Ionic compounds: Will not dissolve in \_\_ Do not consist of \_\_\_\_\_ \_\_\_\_\_ structure As a solid, are \_\_\_\_\_\_ When molten (liquid) & aqueous, they are \_\_\_\_ \_\_\_
Oil Molecules Crystalline Insulators Good conductors
Bonding:
Elements with __ electronegativity will pull electrons away from those with __ electronegativity
High
Low
Bonding:
Cation and anion get pulled apart
Dissociate
Bonding: Tend to form ions with what charge: Group 1: Group 2: Group 15: Group 16: Group 17:
\+1 \+2 -3 -2 -1
Bonding:
The ___ of covalent compounds tell you how many ____ you have of each element
Prefixes
Atoms
Bonding:
A positively charged ion
Cation
Bonding:
If you have more than 1 polyatomic ion, you must put the _____ outside of ______
Subscript
Parentheses
Bonding:
Forms when 2 or more ions combine
Repeating pattern of ions
Transfer electrons
Ionic compound
Bonding:
A separation of positive and negative charge
Dipole
Bonding:
Since d&f block elements have lots of orbitals with nearly overlapping energies, many form ions with more than one _____
Charge
Bonding:
The difference in electronegativity between bonding atoms
Ionic character
Bonding: If binary (2 elements): "ide"-> hydro-----ic acid If tertiary (polyatomic ion): "ates" -> -----ic acid. "ites"-> ------ous acid
Rules for naming acids
Bonding:
If first element is a metal it’s _____
If first element is a nonmetal, its ____
If first element is hydrogen, its ____
Ionic
Covalent
Acid
Bonding:
Related to polarity
The greater the difference in _____, the greater the ___ ____
Ionic > Polar covalent > non polar covalent
Bond strength
Electronegativity
Bond strength
Bonding:
____ bonds form when electrons are transferred from 1 atom to another
____ bonds form when electrons are shared between 2 atoms
____ bonds form when electrons between atoms are delocalized and are free to move around in an electron sea (this is what permits electrical conductivity)
Ionic
Covalent
Metallic
Bonding:
Since an s-orbital can hold _ electrons and a p-orbital can hold _, atoms are stable when they have _ electrons in their valence shell
2
6
8
Bonding:
Anions form when atoms ___ electrons, leaving more ____ than ____
Most ____ will form anions
Gain
Electrons
Protons
Nonmetals
Bonding:
____ have no special name
Na+ is the _____ ___
Cations
Sodium ion
Bonding:
Are nonreactive because they have full s & p orbitals
Noble gases
Bonding: There are 3 categories of bond types: 1)\_\_\_\_: NaCl 2)\_\_\_\_: O2 3)\_\_\_\_: Al, Ca, Fe
Ionic
Covalent
Metallic
Bonding:
The ____ ion
Carbon has _ valence electrons and Nitrogen has _
A ____ bond forms between C & N sharing _ electrons each
Carbon ___ an extra electron
The whole molecule acquires a ___ charge
Cyanide 4, 5 Covalent 3 Steals Negative
Bonding:
A negatively charged ion
Anion
Bonding:
Orbitals are stabilized when they are completely __, completely ___, or to a lesser extent, ______
Full
Empty
Half-full
Bonding:
Covalent compounds:
Have __ melting points
Can be ____, ____, or ____ at room temperature
They may dissolve in ___ or ___, but usually not both
Will not ____ in water
Will not ____ ____ (few exceptions)
Low Solid, liquid, gas Water, oil Dissociate Conduct electricity
Bonding:
Ionic Character:
If the difference is less than or equal to 0.4, bond is _____.
Electrons are ____ ____.
If the difference is greater than 0.4 and less than or equal to 1.6, bond is ____ ____.
Electrons are ____ ___.
If the difference is greater than 2, bond is ___.
Electrons are ____.
Covalent Equally shared Polar covalent Unequally shared Ionic Transferred
Bonding:
Name cation, then the anion
If cation charge fixed, use name from periodic table
If cation charge isn’t fixed, use Roman numeral in parenthesis to specify charge of cation
If anion is single element, drop last syllable and add “-ide”
If anion is a polyatomic ion, use that name
Rules for naming ionic compounds
Molecular shapes:
Electron pairs on central atom:
Shared: 6
Unshared: 0
sp^3d^2
Octahedral
Molecular shapes: To determine the direction of the dipole moment: 1) 2) 3)
1) determine the types of bonds within the molecule
2) draw individual bond dipoles for polar bonds
3) check for molecular symmetry
Molecular shapes:
In order for a molecule to be polar, it must:
1)
2)
1) Have polar bonds
2) Be asymmetrical
Molecular shapes:
Unequal forces between bonds, polar
Asymmetrical
Molecular shapes:
Electron pairs on central atom:
Shared: 3
Unshared: 0
sp^2
Trigonal Planar
Molecular shapes:
Electron pairs on central atom:
Shared: 2
Unshared: 0
sp
Linear
Molecular shapes:
Electron pairs on central atom:
Shared: 2
Unshared: 3
sp^3d
Linear
Molecular shapes:
1) If H is having FON, what IMFAs does the molecule have?
2) If H is not having FON, what IMFAs does the molecule have?
1) hydrogen bonding, dipole-dipole, LDFs
2) dipole-dipole, LDFs
Molecular shapes:
What are the 5 IMFAs?
1) Ion-ion
2) Ion-dipole
3) Hydrogen bonding
4) Dipole-dipole
5) London Dispersion Forces (LDFs)
(Only 3, 4, & 5 on quiz)
Molecular shapes:
Electron pairs on central atom:
Shared: 2
Unshared: 2
sp^3
Bent
Molecular shapes:
Electron pairs on central atom:
Shared: 4
Unshared: 2
sp^3d^2
Square Planar
Molecular shapes:
Electron pairs on central atom:
Shared: 2
Unshared: 1
sp^2
Bent
Molecular shapes:
Electron pairs on central atom:
Shared: 5
Unshared: 0
sp^3d
Trigonal Bipyramid
Molecular shapes:
What IMFA do all molecules have?
London dispersion forces
Molecular shapes:
Electron pairs on central atom:
Shared: 3
Unshared: 2
sp^3d
T-shaped
Molecular shapes:
Electron pairs on central atom:
Shared: 4
Unshared: 1
sp^3d
See-saw
Molecular shapes:
Equal forces from all sides, nonpolar
Symmetrical
Molecular shapes:
Electron pairs on central atom:
Shared: 5
Unshared: 1
sp^3d^2
Square Pyramid
Molecular shapes:
What shapes are always non polar because they are symmetrical?
(Only when central atom is surrounded by the same elements)
(Put shared and unshared and shape)
2,0: linear 3,0: trigonal planar 4,0: tetrahedral 5,0: trigonal Bipyramid 6,0: octahedral 2,3: linear 4,2: square planar
Molecular shapes:
Electron pairs on central atom:
Shared: 3
Unshared: 1
sp^3
Trigonal pyramid
Molecular shapes:
In a polar bond, the higher electronegative element is slightly ____ and the other is _____
Negative
Positive
Molecular shapes:
Electron pairs on central atom:
Shared: 4
Unshared: 0
sp^3
Tetrahedral
Uncertainty in Measurement:
Use ___ ___ to eliminate all placeholding zeroes
Scientific notation
Uncertainty in Measurement:
When a number is in scientific notation, the number being multiplied by 10* has all ____ ____
Ex: 2.590 x 10^4 has _ sig figs
Significant figures
4
Uncertainty in Measurement: Sig fig rules do not apply when: 1) not \_\_\_\_\_ something 2) \_\_\_\_\_ something 3) have a defined \_\_\_\_ or \_\_\_\_
Measuring
Counting
Values
Counts
Uncertainty in Measurement:
Even odd rule: if a number ends in exactly .5, look at the number in front of it. If the number in front is even, round ___. If the number in front is odd, round ___.
Down
Up
Uncertainty in Measurement:
How closely several measurements agree
Precision
Uncertainty in Measurement:
Rules for multiplying/dividing sig figs:
The answer cannot have more sig figs than there are in the measurement with the ___ number of sig figs
Least
Uncertainty in Measurement:
How close a measurement is to the true value
Accuracy
Uncertainty in Measurement:
To identify how many sig figs are in a measurement, follow these rules:
1) All _____ digits are significant
2) All ___ between nonzero digits are always significant
3) ___ in front of nonzero digits are never significant
4) ____ after all nonzero digits are only significant if a ___ __ is present anywhere in the number
Nonzero
Zeroes
Zeroes
Zeroes, decimal point
Uncertainty in Measurement:
Rules for adding/subtracting sig figs:
The answer cannot have more __ to the __ of the decimal point than there are in the measurement with the ____ number of ___ to the right of the decimal point
Digits
Right
Smallest
Digits
Uncertainty in Measurement:
Measurements are made __ place beyond where the measuring device is marked. Includes all certain digits and _ estimated digit
1
1
The Mole:
Formula that shows the composition of a compound in terms of relative number and kinds of atoms in the simplest, whole number ratio
Empirical formula
The Mole:
Masses of _ are so small that expressing them in terms of _ is not convenient
Atoms
Grams
The Mole:
Mass in grams of 1 mole of an element. The unit is /. It is numerically equivalent to the __ __ of an element
Molar mass
Grams/mole (g/mol)
Atomic mass
The Mole:
1 mole of a substance = ___ x __^__ particles
6.02x10^23
The Mole:
Mole: Number of atoms in exactly _ grams of __-__. It is the SI base unit for _.
12
Carbon-12
Quantity
The Mole:
To convert between moles and mass, we use the __ __
Molar mass
The Mole:
Percentage by mass of each element in a compound. Helps verify a substance’s identity
Percent composition
The Mole:
The actual formula of a single molecule of a compound. Sometimes can be same as empirical formula
Molecular formula
The Mole:
The number of particles present in 1 mole of a substance is called ___ ___
Avogadro’s number
Chemical Equations:
Molecules that act to lower the activation energy of a reaction. Are not consumed by the reaction, but play a role in enabling the reaction to proceed with __ energy requirements
Catalysts
Lower
Chemical Equations:
Change in __ is always a physical change. For proof, you need a __ __ to show at least one new substance has formed
State
Chemical analysis
Chemical Equations:
A measure of the amount of molecules per set volume
Concentration
Chemical Equations:
Compound is completely oxidized and broken up by oxygen (w/ flame). CO2 & H2O are always the products
Combustion (or burning)
Chemical Equations:
Direct combination of elements to produce a compound. Single product
Synthesis (or combination) reaction
Chemical Equations:
__ reactions absorb energy, so are only spontaneous if the ___ is higher than the ___ ___
Endothermic
Temperature
Activation energy
Chemical Equations:
Substances created
Products
Chemical Equations:
Process by which one or more substances change to produce one or more different substances
Chemical reactions
Chemical Equations:
Compounds break down into constituent elements and/or smaller molecules. 1 reactant
Decomposition (or analysis) reaction
Chemical Equations:
The higher the ___ the faster the reaction
Concentration
Chemical Equations:
In order for a reaction to occur the reactants must:
1) ____
2) have enough __ __
3) have the correct ___ when they collide
4) form ____ favorable products
Collide
Activation energy
Orientation
Energetically
Chemical Equations:
Amount of energy needed to initiate a reaction. Without enough __, molecules that react very violently will __ __ at all
Activation energy
Energy
Not react
Chemical Equations:
Original substances
Reactants
Chemical Equations:
The more __ present in a confined space, the greater the chance they will __
Molecules
Collide
Chemical Equations:
Takes a long time for molecules in the _ of “chunk” to contact other reactant (have _ surface area)
Center
Low
Chemical Equations:
To speed up reaction, break __ pieces into __ pieces with high surface areas
Big
Smaller
Chemical Equations:
Balance equations by inserting __ in front of the compounds
Coefficients
Chemical Equations:
An ion from one compound switches places with similarly charged ion from another compound
Double replacement (or displacement)
Chemical Equations:
For a reaction that proceeds at room temperature, ___ the temperature will speed up the reaction. There are exceptions
Increasing
Chemical Equations:
Once started, ___ reactions often produce enough energy to maintain a ___ reaction
Exothermic
Spontaneous
Chemical Equations:
An element will replace a similar element in a compound
Single Replacement (or displacement)
Chemical Equations:
A representation of a chemical reaction that uses symbols to show the relationship between the reactants and the products
Chemical equations
Chemical Equations:
Reactions that occur spontaneously at room temperature have __ activation energies. Reactions that require high temperatures have __ __ activation energies
Low
Very high
Chemical Equations: Evidence of a reaction 1) 2) 3) 4) 5)
1) changes in energy
2) formation of a gas
3) formation of a precipitate
4) change in color
5) odor
Chemical Equations:
Affects the number of molecules that have the correct activation energy for the reaction
Temperature
Chemical Equations:
How to write net ionic equations:
1) Write all the compounds that are soluble as ions (put charges)
2) cross out any spectator ions that are on both the left and right side
Chemical Equations: Physical state of substances (s): (l): (g): (aq):
(s) : solid
(l) : liquid
(g) : gas
(aq) : aqueous solution (dissolved in water)
Chemical Equations:
What drives a reaction?
Makes a gas (which escapes)
Makes a solid (which precipitates)
Makes water
Changes oxidation state (charge)
Chemical Equations:
A measure of the ratio of the amount of surface to volume of a substance
Surface area
Chemical Equations:
Triangle over arrow indicates __ added to the reaction.
Double arrow indicates an ____ reaction
A substance’s name over the arrow indicates the use of a ___
Heat
Equilibrium
Catalyst
Chemical Equations:
Strong acids and strong bases combine to form water and a salt. Acids: have an _ in front
Bases: usually __
Neutralization (or acid-base) reaction
H
OH
