Chemistry Exam 1 Flashcards
Chapters 1-5
Compounds
A substance composed of two or more different chemical elements that are chemically bonded together in a fixed ratio
Mixtures
A substance made up of two or more different components that are physically combined but not chemically bonded together
Element
A pure substance that cannot be broken down into simpler substances by chemical means
Pure Substance
A material that has a constant composition and distinct properties throughout, with no variation in its chemical make-up
Homogenous Mixture
The components are uniformly distributed with no visible boundaries between them. Examples: air, salt water, vinegar
Heterogenous Mixture
The components are not uniformly distributed and have visible boundaries between different phases/components. Examples: sand and water, salt and oil
Changes In State: Which ones require an input of energy?
Melting, Boiling, Sublimination (solid to gas)
Changes in State: Which ones release energy?
Condensation, Freezing, Deposition (gas to solid)
Chemical vs Physical Changes
Chemical changes alter the chemical composition (tarnishing of silver), while physical changes just change in color, boiling point, etc
Steps of Scientific Method
Step 1: Make Observations
Step 2: Develop a Hypothesis
Step 3: Test the hypothesis through experiments
Step 4: Develop a law (summarize consistent outcomes)
Step 5: Develop a theory (explanation of law)
Scientific Notation
- Move the decimal point so there is only one non-zero digit to the left of the decimal
- Count the number of places the decimal moved
- The digit(s) left is the coefficient
- The number of places the decimal moved becomes the exponent of 10
Sig Fig Rules
- Answer can’t have more sig figs than lowest # sig figs in the calculation
- Leading zeros are not significant UNLESS marked by decimal (3500.)
- Trailing zeros are signif. (.3000)
Density equation
D=M/V
Specific Heat Capacity
Amount of energy needed to raise a certain mass of a substance by 1 degree Celsius
Exothermic Reactions
Release Energy
Endothermic Reactions
Absorb energy
Law of Conservation of Matter (Dalton)
When a reaction takes place, matter is neither created nor destroyed
Percent by Mass
mass of oxygen in compound divided by total mass of compound times 100
Dalton’s Atomic Theory
- All matter is made up of atoms
- Atoms can neither be created or destroyed
- Atoms of a particular element are alike
- Atoms of different elements are different
- A chemical reaction involves the union or separation of individual atoms
Rutherford’s Gold Foil Experiment
- Fired alpha particles at metal foil
- Expected a straight flight path for the particles
- Some particles flew through and some deflected at a wide angle or “bounced backward”
- made him realize there must be something dense inside the atom and negative space
Rutherford’s Model of the Atom
- An atom is mostly negative space
- A nucleus is a tiny, massively dense, positive-charged unit in the atom (contains protons and neutrons)
Atomic Number
- # of protons in the nucleus
- Determines identity of the atom
- An atom with one proton is always a hydrogen
Mass Number
Number of protons plus neutrons
Isotopes
Atoms with the same number of protons but differing numbers of neutrons (exhibit same chemical properties still)
Atomic Mass
The actual mass of an atom (given in AMUs or Daltons)
Law of Octaves
Elements that are 8 elements apart by mass react in a similar manner (aka chemical periodicity)
Law of Mendelev
Properties of elements recur in regular cycles (periodically) when elements are arranged in increasing atomic mass
Periods on the Periodic Table
-The horizontal rows on the table
-Numbered 1-7 down the table
Groups on Periodic Table
-The verticle columns on the table
-Also called families
-Either numbered in Roman numerals or Arabic numbers
Metals
-Shiny solids
-Bendable, malleable
-Conductors of heat and electricity
Nonmetals
-Brittle
-Do not conduct electricity or heat well
Metalloids
-Can act as a metal or nonmetal
Bohr’s Theory
-Electrons are attracted to the nucleus because of their charge
-Energy must be added for the electron to move away from the nucleus
-When the electrons move back to their original positions, energy is released
Electrons in Orbits
-The closer electrons are to the nucleus, the less energy they possess
Valence Shell
-Outermost shell of the Bohr model
-The # of electrons in valence shell determines the chemical properties
Ground State
Lowest energy state (electrons as close to the nucleus as possible)
Excited State
Any state higher in energy than ground state (1 or more electrons travelled away from nucleus)
Subshells & Electron Configurations
- n=1 s2
- n=2 s2 p6
- n=3 s2 p6 d10
- n=4 s2 p6 d10 f14
*s ALWAYS comes before the next d
What Shell Has the Highest Energy?
F subshell has highest, lowest is s subshell
Electron Configuration Example
1s2 2s2 2p6 3p2 3p6 4s2 3d10
How to Find Valence Electrons
-Do electron configuration (last shell shows valence electrons - but you must add s and d shells together if most outer)
-Find element valency and multiply that and add to any other elements
-count of lewis structure
Atomic Size and Periodic Table
There’s a decrease in atomic size when you move left to right across a period (increase in atomic size pulls electrons inwards, resulting in decreasing atomic size)
Ionization Energy and Periodic Table
-Moving from left to right, the ionization energy increases across a period
-Moving down a family, ionization energy decreases
Ionization Energy
The energy required to remove the outermost electron from an atom
Element Size and Periodic Table
-Moving across a period, size decreases
-Moving down a group, size increases
*smaller size= more attraction between nucleus and electrons
Alkali metals
-Group 1A, have low ionization energy and give up electrons easily
Halogens
-Group 7A, have high ionization energy and would rather gain electrons
Alkaline-Earth Metals
-Group 2A, have low ionization energy and give up electrons
Noble Gases
-Group 8A, have a full outer shell and do not form compounds easily
Cation
-An ion with a positive charge
-Forms when an atom loses one or more electrons, leaving it with less electrons than protons
Anion
-An ion with a negative charge
-Forms when an atom gains one or more electrons, leaving it with more electrons than protons
Do Metals or Nonmetals tend to gain electrons?
-Metals tend to lose electrons
-Nonmetals tend to gain electrons
Ionic Compounds
-Formed by an attraction of ions with opposite charges
-Formed by metals and nonmetals
Electrostatic Force
The smaller the distance and the larger the charges, the stronger the force holding oppositely charged ions together
Molecules
-A collection of atoms bound together without forming ions (so they’re electrically neutral
-Have ionic or covalent bonds
-Can’t be broken down physically but can chemically
Molecular Compound
-Composed of 2 or more types of elements
-Examples: CH4 and H2O
Elemental Substance
-Composed of only one type of element
-Examples: O2, and Cl2
Covalent Bond
-Bond that occurs between electrons in one atom and the nucleus of another
-Electrons are usually shared between atoms
Ionic Bond
-Formed between metals and nonmetals: Metals tend to lose electrons and become cations, while nonmetals gain those electrons to become anions.
-Electrostatic attraction: The opposite charges of the cations and anions attract each other to form the ionic bond
Core Electrons
Almost never involved in bonding because of their proximity to the nucleus
Valence Electrons
-Electrons that are shared between neighboring atoms
-Responsible for covalent bonding
Energy in Bonds
-Bond formation ALWAYS releases energy
-Bond breaking ALWAYS takes energy
Diatomic Molecule
A molecule composed of only two atoms of the same or different chemical elements (O2, Cl2, H2)
Electronegativity (EN)
-Measurements for how much an atom wants another electron
-Determines if electrons are transferred or shared when a bond is formed
-Greater the EN, the more likely electrons will be drawn to an atom
Pure Covalent Bonding
Result when 2 atoms have identical electronegativity
Covalency
-Non-polar: 0-0.4
-Polar: 0.4-1.7
-Ionic 1.7-4.0
* covalency increases left to right and down to up
Polar Covalent Bonding
Result when atoms have differing electronegativity (unequal electron sharing)
Oxidation
-Removal of hydrogen or electrons
-Addition of oxygen
Reduction
-Addition of hydrogen or electrons
-Removal of oxygen
