CHEM111 CHAPTER 8-12 Flashcards
1
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Thermodynamics
A
Study of energy and temperature changes
2
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Potential energy
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Stored energy that matter possesses due to its position, condition and/or composition (cohesive force)
3
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Kinetic energy
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Energy of motion (disruptive force)
4
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What happens when a substance absorbs heat energy
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The particles within the substance vibrate and move faster
5
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Equation that represents energy, heat and work
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ΔE=q+w. Triangle E represents change in energy, q is heat, and w is work
6
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Exothermic change
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Releases heat energy (-q)
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Endothermic change
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Absorbs energy (+q)
8
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Law of conservation of energy
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Also called the first law of thermodynamics that states energy cannot be created or destroyed (ΔEsystem)=-(ΔEsurroundings)
9
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Specific heat
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Relationship between heat, mass, and temperature. It is the amount of heat required to raise the temperature of one gram of material by one degree Celsius. s=q/mΔT
10
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Coffee cup calorimetry
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When water is placed in an insulated cup covered with a cork lid and a thermometer. The system to be studied is placed inside the coffee cup and the reaction or change is allowed to take place
11
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Bomb calorimetry
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Measures the heats of reactions for fuels and other high energy chemical changes. T substance is placed in a heavy steal container and outside of the container is a second container filled with water and equipped with a thermometer
12
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Extensive property
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Energy absorbed or released in a reaction
13
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Fuel value
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The amount of heat energy that can be released by a combustion reaction of a certain substance (expressed in terms of energy per unit gram)
14
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Reaction enthalpy
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Amount of heat energy that is absorbed or released in a chemical reaction at constant pressure represented by triangle Hrxn
15
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Physical state depends on (4):
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- Chemical identity of the matter (what it is)
- Temperature of the matter
- Pressure the matter is subjected to
- The intermolecular forces that hold it together
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Q
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Representation of total energy released or absorbed by a system
17
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Specific heat capacity
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When heat capacity is given per gram of substance
18
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Molar heat capacity
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When heat capacity is given per mole of substance
19
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Sublimination
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When a material goes from solid to gas without going through the liquid state (endothermic - energy required to break intermolecular bonds)
20
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Deposition
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When a material goes from gas to solid without going through the liquid state (exothermic - energy released as intermolecular bonds form)
21
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Ionic bonds
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Attraction between oppositely charged ions
22
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Covalent bonds
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Attraction between two nonmetals and results when two atoms share electrons
23
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Octet rule
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States atoms are stabilized by 8 electrons in their valence level
24
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Expanding octet
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A bonding arrangement in which an atom has 10 or 12 valence electrons. This is ONLY possible with elements in rows 3-7 on periodic table
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Single covalent bond
Two atoms share one pair of valence electrons
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Double covalent bond
Two atoms share two pairs of valence electrons
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Triple covalent bond
Two atoms share three pairs of electrons
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Coordinate covalent bond
Both electrons of a shared pair originate from the same atom
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Strength of bond
The more electron pairs being shared, the greater the strength of bond. increase as it goes single -> double -> triple
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What element is never the central atom
Hydrogen because only two valence electrons. Or Fluorine
31
Localized electron model
Assumes all electrons are in pairs thus it does not handle odd-electron molecules very easily
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Steps to drawing Lewis Structures (4)
- Add up all the valence electrons
- Frame the structure using single bonds (typically the atom that is nearer to the lower left of the periodic table is the central atom)
- Fill octets of the outer atom first and add lone pairs
- Fill the octet on the central atom
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Formal charges
the charge that would reside on the atom if all of the bonding electrons were shared equally.
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Resonance structures
Set of Lewis structures that show how electrons are distributed around a molecule or ion. Resonance structures are used when a single Lewis structure cannot adequately depict the structure (for 2nd or 3rd covalent bonds, and lone pairs)
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Valence shell electron pair repulsion (VSEPR)
Model based on the idea that electrons repel each other, and therefore they occupy regions that are as far away from each other as possible
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Electronic geometry
Arrangement of electrons around a central atom
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Molecular geometry
Arrangement of atoms within a molecule
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Two VSEPR electron pairs
linear
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Three VSEPR electron pairs
Trigonal planar (120 deg) - Bent
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Four VSEPR electron pairs
Tetrahedral (109.5 deg) - Trigonal pyramidal
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Five VSEPR electron pairs
Trigonal Bipyramidal
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Six VSEPR electron pairs
Octahedral
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What shape is a water molecule
Tetrahedral
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Electronegativity
Measure of the ability of that atom to attract shared electrons towards itself and is affected by atom size, nuclear charge, and the number of non valence electrons
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Pauling scale
Numerical scale of electronegativities based on bond-energy calculations for different elements joined by covalent bonds
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Pure covalent bond
Bond between two identical atoms or atoms of the same electronegativity (equal sharing of electron)
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Polar covalent bond
Covalent bond between atoms of intermediate difference in electronegativity (unequal sharing of electrons)
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Greater difference in electronegativity = ____
The greater the polarity of the bond
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Bond polarity
Measure of the inequality in sharing of electrons in a bond
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Dipole
An overall polarity in which different sides of a molecule have slight positive and negative charges
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Which element is the most electronegative
Fluorine. It has the strongest pull for electrons
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Which elements are the least electronegative
Cesium and francium. It has the weakest hold on electrons
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Bond summary (3):
- Covalent: Difference in electronegativity less than 0.5
- Polar covalent: Difference between 0.5 and 2.0
- Ionic: Difference greater than 2.0
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Types of intermolecular forces (3):
- Hydrogen bonding
- Dipole-dipole forces (van der waals forces)
- London dispersion forces (van der waals forces)
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Dipole-dipole forces
Attractive forces between the positive end of one polar molecule and the negative end of another polar molecule. Can interact with each other electrostatically (positive end of one molecule lines up with negative end)
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Hydrogen bonding
Strong dipole-dipole forces when H is attached to a highly electronegative elements such as N, O, or F (very polar bond)
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London dispersion forces
Results from the motion of electrons that creates temporary dipoles in molecules.
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Strength of intermolecular forces least to greatest
London dispersion --> dipole-dipole --> hydrogen bonding
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Key pointers of hydrogen bonding (2):
- Requires donor and acceptors (donor is H bound to an N, O, F, and acceptor is a lone pair on an N, O, F)
- Intermolecular hydrogen bonding also seen in alcohols (organic molecule containing O-H group) and amines (organic molecule containing N-H group)
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Intermolecular forces
Usually [weak] forces of attraction or repulsion between individual molecules
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Strength of London dispersion also depends on what factors
Depends on molecular shape/size. Electrons in elongated molecules are more easily displaced than those in small/symmetrical molecules.
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N-pentane
Large contact area, strong attraction
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Isopentane
Less surface area, less attraction
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Neopentane
Small contact area, weakest attraction
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Viscosity
Measure of a liquids resistance to flow. Liquids with strong intermolecular forces are highly viscous
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Types of crystalline solids (6):
- Ionic solids
- Metallic solids
- Molecular solids
- Network atomic solids
- Polymer network
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Ionic solids
Extended network of ions held together by ion-ion interactions (NaCl, MgO)
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Metallic solids
Extended network of atoms held together by metallic bonding (Cu, Fe)
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Molecular solids
Discrete molecules held together by intermolecular forces (HBr, H2O)
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Network atomic solids
Extended network of atoms held together by covalent bonds (C, Si)
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Polymer network
Macromolecule where each constitutional unit is connected to each other's constitutional unit
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Ideal gas law
Relationship between pressure, volume, and temperature, and the number of moles of an ideal gas expressed as PV = nRT where R is the gas constant
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Barometer
Device used to measure atmospheric pressure. A long tube is filled with a liquid (usually mercury) then turned upside down in a liquid. Gravity pulls the liquid down, creating a vacuum in the top of tube and the pressure from the outside air pushes the liquid up the tube. The higher the air pressure, the higher the liquid is pushed into the tube.
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Gauge pressure
Difference between a compressed gas pressure and atmospheric pressure
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Atmosphere (atm)
Unit of gas pressure (1atm=760mmHg)
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The gas laws
Generalizations that summarize in mathematical terms, experimental observations about the relationships among the amount (moles), pressure, temperature and the volume of gas
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Boyles law
As pressure increases, volume decreases. As pressure decreases, volume increases (temperature and number of moles remain unchanged). PV= k
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Charles law
As temperature increases, volume increases. As temperature decreases, volume decreases (pressure and number of moles remain unchanged) V/T = b
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Gay-Lussacs law
As temperature increases, pressure increases. As temperature decreases, pressure decreases (volume and number of moles remain unchanged) P/T =c
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Combined gas law equation
P1V1/T1 = P2V2/T2
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Avogadro's law
States that, if pressure and temperature are constant, the volume of a gas is proportional to the number of moles of gas present. V/n =a
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Partial pressure
Pressure caused by one gas in a mixture. Adding up all the partial pressures gives the total pressure
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Universal gas constant (R)
R= 0.08206 L atm/K mol
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Effusion
Process of a gas escaping from a container; lighter gases effuse more quickly than heavier gases
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Diffusion
Term used to describe the mixing of gases; Particles randomly move, and as they do, they slowly spread from areas of higher concentration to lower concentration
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How does the pressure change if the temperature of the gas increases?
At higher temperatures, the particles move faster
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What happens to the pressure of a gas if the volume increases?
If the volume of a container increases, the particles have to travel farther before reaching the walls of the container. As a result, the particles collide with the container less frequently
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What happens to the pressure of a gas if the number of particles in a container is increased?:
More gas particles mean the particles will strike the sides of the container more often, creating more pressure (think of pumping a tire)
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Solution
A homogenous mixture of 2 or more substances (can be gases, liquids, solids)
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Solute
Substance that is dissolved in a solution; smaller component
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Solvent
A major component of a solution
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Solvent
The component of the mixture present in the greatest amount
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Aqueous solution
When water is the solvent
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Dilute
A solution that has very little solute
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Concentrated
A solution that has lots of solute
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Saturated
A solution containing the maximum amount of dissolved solute
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Percent by mass equation
mass of solute/mass of solution x 100%. To find measure of concentration
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Percent by volume equation
Volume of solute/volume of solution x 100%. To find measure of concentration
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Molarity (M)
Moles of solute/litre of solution; measure of concentration
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Molality (m)
Moles of solute/kg of solvent
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Percent composition
Percent of the solute present in the solution
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Parts per million (ppm)
A measure of concentration used for dilute solutions (1g of solute per 1,000,000 grams of solution); mass of solute/mass of solution x 10^6
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Parts per billion (ppb):
A measure of concentration used for dilute solutions (1g of solute per 1,000,000,000 grams of solution); mass of solute/mass of solution x 10^9
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What do colligative properties depend on
Depend only on the number of molecules or ions of the solute that are dissolved. Examples of colligative properties are boiling points, freezing points, etc
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Osmosis
Tendency of water to move towards regions of greater concentrations
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Freezing point depression
A colligative property of water; presence of solute in an aqueous solution lowers the freezing point below that of pure water
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Boiling point elevation
A colligative property of water; the presence of solute in an aqueous solution raises the boiling point above that of pure water
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Semipermeable membrane
Lipid bilayer where water can pass back and forth through membrane, but most molecules and ions cannot
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Precipitation reactions
When two solutions are combined and a solid product is formed
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Metal displacement reaction
A reaction between two metals in which one metal is oxidized to its ionic form while the other metal ion is reduced to its elemental form
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Gravimetric analysis:
A technique that uses the mass of a precipitate to determine the concentration of a reactant
112
Acids (5):
- ACIDS ARE PROTON (H+) DONORS AND PRODUCES H+ OR H3O+ IN WATER
- Sour taste
- Turns litmus red
-Some metals dissolve in acid, giving off H2 gas
- Examples: Vinegar, citric acid
113
Base (6):
- BASES ARE PROTON (H+) ACCEPTORS AND PRODUCES OH- IN WATER
- Bitter taste
-Turns litmus blue
- Fatty substances dissolve in base
- Slippery to touch
- Examples: Soaps, cleaners containing ammonia
114
What is the number one produced chemical worldwide?
Sulfuric Acid
115
Arrhenius Acid
A hydrogen containing compound that dissolves in water releasing H+ ions and are polar covalent molecules when not in solutions but form ions in presence of water
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Arrhenius Base
A hydroxide containing compound that dissolves in water to give OH- ions and are ionic compounds that dissociate in water
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Arrhenius theory
Theory of ionization. States that acids are hydrogen containing compounds which give H+ ions or protons in dissociation in water and bases are the hydroxide compounds which give OH- ions in dissociation in water
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Bronsted-Lowry Theory (more common)
States that a substance can function as an acid only in the presence of a base; a substance can only function as a base in the presence of an acid
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Ionization
When molecular compounds dissolve in water to produce ions
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Dissociation
When ionic compounds dissolve to produce ions and occurs when water molecules "pull apart" the ionic crystal
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Amphoteric substance
Substances that can act as an acid or base such as water
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Hydronium ion
H3O+, the ion formed in acidic aqueous solutions
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Protonation
Gain of an H+
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Deprotonation
Loss of H+
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Monoprotic acids
Acids that can release only one proton into solution
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Polyprotic acids
Acids that donate two or more protons
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Diprotic acids
Acids that can release 2 protons (Sulfuric acid, carbonic acid)
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Tripotic acid
Acids that can release 3 protons (Phosphoric acid)
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Strong acids
Acids that completely ionize in water
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Weak acids
Acids that only partially ionize in water (eg. HF)
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Equilibrium
Two opposite processes taking place at equal rates
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Conjugate acid
Acid on the right hand side of chemical equation in an acid-base equilibrium; acid formed when a base reacts with H+
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Conjugate base
Base on the right hand side of a chemical equation in an acid-base equilibrium; base formed when an acid releases an H+
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Neutralization reaction
Reaction in which an acid and a base combine; acids and hydroxide bases continue to produce water and an ionic compound = salt
135
Acid rain
The effect observed when nonmetal oxides combine with moisture in the atmosphere to produce acidic rainfall
136
Common strong acids (7):
- Nitric acid
- Sulfuric acid
- Perchloric acid
- Chloric acid
- Hydrochloric acid
- Hydrobromic acid
- Hydroiodic acid
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Strong electrolyte
Soluble salts, strong acids and bases
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Weak electrolyte
Weak acids and bases
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Non-electrolyte
Molecular compounds
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Acid-base titration
Determination of an unknown concentration of an acid or base by exactly neutralizing it with an acid or base of known concentration
141
Litmus
Turns blue in the presence of a base and red in the presence of acid
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Phenolphthalein
Turns bright pink in a basic solution but colourless in an acidic solution
143
pH Paper
Paper containing a blend of indicators that can be used to estimate pH based on color; produces colors ranging from deep red to deep blue
144
pH scale
Scale that indicates the relative concentration of acid or base in an aqueous solution; pH is defined as the negative log of the hydronium concentration.
145
pH scale indication
7 is neutral, lower pH is acidic, higher pH is basic
146
Buffer
A solution that can resist pH change upon addition of an acidic or basic components and able to neutralize small amounts of acid or base thus maintaining the pH of the solution relatively stable
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How are buffers able to function
Buffers are able to resist pH change because the conjugate acid and conjugate base are both present in appreciable amounts of equilibrium
148
Buffers contain either....
- A mixture of a weak acid and its conjugate base
- A mixture of a weak base and its conjugate acid
149
pH of fluid inside cells
6.8
150
pH of bloodstream
7.4
