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SAT II: Chemistry > Vocabulary > Flashcards

Vocabulary Flashcards

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1
Q

Law of Definite Composition/Proportions

A

compounds are made by combining atoms of two or more elements in a definite proportion or ratio by mass

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2
Q

physical properties

A

observed with senses; physical state, color, odor, solubility in water, density, melting point, taste, boiling point, hardness

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3
Q

chemical properties

A

can be observed in regard to whether or not a substance changes chemically; result of reacting with other substances; Ex: iron rusting, nitrogen not burning, gold not rusting, sodium and water react, silver and water don’t react, water can be decomposed by electric current

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4
Q

Law of Conservation of Mass

A

mass of reactants equals mass of products

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5
Q

Law of Conservation of Energy

A

energy is neither gained nor lost in physical or chemical changes

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6
Q

Law of Conservation of Mass and Energy

A

mass and energy are interchangeable under special conditions

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7
Q

Kelvin and Celsius

A 
K = C + 273
C = K - 273
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8
Q

Joule and calorie

A

1 calorie = 4.18 joules

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9
Q

accuracy

A

how close you have come to the true value

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10
Q

precision

A

how repeatable the results are

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11
Q

+/- Sig figs

A

to the least common place

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12
Q

Mult or Div Sig Figs

A

least amount

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13
Q

Example of Sig Figs

A
  • all digits except zero
  • zeros between nonzero digits
  • zeros to the right of the decimal point
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14
Q

Dalton’s Atomic Theory

A

Matter = small particles called atoms
atoms of an element have similar weight different than other elements
atoms cannot be subdivided, created, or destroyed
atoms of diff elements combine in simple whole number ratios to form chemical compounds
chemical rections: atoms are combined, separated, or rearranged

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15
Q

J.J. Thompson

A

electron

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16
Q

Rutherford

A

protons

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17
Q

James Chadwick

A

neutrons

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18
Q

Niels Bohr

A

orbitals/energy levels

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19
Q

Henry Moseley

A

atomic numbers; periodic law: properties of elements are periodic functions of their atomic numbers

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20
Q

isotopes

A

same number of protons; diff masses bc of neutrons

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21
Q

Heisenberg uncertainty principle

A

it’s impossible to know both the precise location and precise velocity of a subatomic

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22
Q

Pauli Exclusion Principle

A

No two electrons can have the same four quantum numbers

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23
Q

Aufbau Principle

A

an electron occupies the lowest energy orbital that can receive it

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24
Q

Hund’s Rule of Maximum Multiplicity

A

pairing will occur with the addition of one more electron to each orbital

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25
Q

Periodic Table Properties

A

acid properties increase ——>

ionization energy increases ——>
nonmetallic props increase —–>
electronegativity —–>

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26
Q

radioactive emissions

A
  • alpha ray: +
  • beta ray: - and high speed
  • gamma ray: no charge
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27
Q

half-life

A

time required for half of the atoms or a radioactive nuclide to decay; carbon-14 has a half life of about 5,700 years

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28
Q

ionic bonds

A
  • greater than 1.7
  • metal and a nonmetal or polyatomic
  • high melting point
  • won’t conduct electricity
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29
Q

covalent bonds

A
  • diatomic
  • less than .4 is non polar
  • greater than .4 is polar
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30
Q

metallic bonds

A
  • “sea of electrons”

- metal and a metal

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31
Q

Intermolecular forces

A

(in order from strong to weak)

  • hydrogen bonds
  • dipole-dipole interaction
  • dipole induced interaction
  • London/Dispersion forces
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32
Q

resonance

A

a hybrid of the possible drawings bc no one lewis structure can represent the situation

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33
Q

VSEPR theory

A

valence shell electron pair repulsion; bond angles

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34
Q

Hybridization theory

A

changes in the orbitals of the valence electrons

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35
Q

Hybrid Arrangements

A
  • linear: 180; two atoms bonded to center
  • bent: 2 atoms bonded and 1 or 2 lone pairs (water 105 degree angle)
  • trigonal planar: 120; 3 atoms bonded to center
  • tetrahedral: 4 atoms bonded to center
  • trigonal pyramidal: 3 atoms bonded to center and 1 lone pair
  • trigonal bipyramidal: 90 and 120; 5 atoms to center
  • octahedral: six electron to center
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36
Q

Hybrid Orbitals

A
  • sp: 180/linear
  • sp2: 120
  • sp3: example is ammonia
  • sp3d2: hybrid of an s and a p electron promoted to d orbitals and transformed into six equal sp3d2 orbitals
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37
Q

Acid Naming Rules

A
  • “ic” acids form “ate” salts
  • “ous” acids form “ite” salts
  • “hydro…ic” acids form “ide” salts
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38
Q

Law of Definite Composition

A

each unit of an element has the same atomic mass and every time the particular compound forms, it forms in the same percent composition

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39
Q

Law of Multiple Proportions

A

some compounds form where the mass of one element stays the same but the mass of the other varies

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40
Q

Activity of Metals

A
  • potassium, calcium, sodium (in cold water)
  • low to high: hydrogen, {tin, iron, zinc, aluminum} in most dilute solutions, magnesium in hot water
  • {sodium, calcium, potassium} in cold water
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41
Q

Standards

A
  • 760 mmHg = 760 torr = 1 atm = 101.325 kPa

- 273 K = 0C

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42
Q

Kinetic Molecular Theory

A
  • matter is composed of tiny particles
  • particles of matter are in constant motion
  • when particles collide, there is no loss of energy
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43
Q

Graham’s Law of Effusion (Diffusion)

A

The rate of a gas is inversely proportional to the square root of its molecular mass

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44
Q

Charles’s Law

A

V₁T₁=V₂T₂ where the pressure is constant and the volume and temp vary directly

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45
Q

Boyle’s Law

A

P₁V₁=P₂V₂ where the temp is constant and the volume and pressure vary inversely

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46
Q

Combined Gas Law

A

P₁V₁/T₁=P₂V₂/T₂

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47
Q

Gay Lussac’s Law

A

P₁/T₁=P₂/T₂ where the volume is constant and the pressure and temp vary directly

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48
Q

Dalton’s Law of Partial Pressures

A

total pressure is equal to the sum of the partial pressures

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49
Q

Converting water -> Mercury

A

-divide the difference by 13.6

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50
Q

Ideal Gas Law

A

PV=nRT (where R can be .0821atm or 8.31kPa)

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51
Q

Ideal Gas Deviations

A
  • least deviations occur at low pressures and high temps

- high deviations occur at high pressures and low temps

52
Q

percent yield

A

actual yield/theoretical yield *100

53
Q

capillary action

A

the attraction of the surface of a liquid to the surface of a solid

54
Q

Le Chatlier’s Principle

A
  • system will react to stress
  • temp: ↑endo or ↓exo
  • pressure: ↑less moles or ↓more moles
  • conc: ↑away or ↓toward
55
Q

Solids

A
  • Crystalline
  • Amorphous
  • Polycrystaline
56
Q

Water’s heat of vaporization

A

40.79 kJ/mol

57
Q

Water’s heat of fusion

A

6.01 kJ/mol

58
Q

anhydrides

A
  • oxides that react with water to from acids and bases
  • metal oxides form bases (known as basic anhydrides or basic oxides)
  • nonmetallic oxides form acids (known as acidic anhydrides or acidic oxides)
59
Q

General Rules of Solubility

A
  • Solids: solubility usually ↑ when temp ↑; pressure has little effect
  • Gasses: solubility usually ↓ when temp ↑; solubility varies in direct proportion to the pressure applied (Henry’s law)
60
Q

Factors that affect Solubility

A
  • ↑ surface area
  • stirring
  • heating
61
Q

% Concentration

A

grams of solute/grams of solution *100 (10⁶ppm or 10⁹ppb)

62
Q

Molarity

A

moles of solute/liters solution

63
Q

Molality

A

moles of solute/kg solvent

64
Q

Mole Fraction

A

moles of interest/total moles

65
Q

dilution

A

M₁V₁=M₂V₂

66
Q

heat of formation

A
  • low: easy decomp

- high: extreme instability

67
Q

heat of combustion

A

(combustion reactions)

68
Q

Conditions for Completion

A

insoluble ppt formed
nonionizing substance formed
gaseous product given off

69
Q

Hydrolysis

A

-salt and water react to form a base (opp of neutralization)

70
Q

Standard State Condition

A

25C and 1atm

71
Q

Hess’s Law of Heat Summation

A

when a reaction can be expressed as the algebraic sum of two or more other reactions; the heat of reaction is the algebraic sum of the heats of these reactions (based on 1st law of thermodynamics)

72
Q

First Law of Thermodynamics

A

the total energy of the universe is constant and cannot be created or destroyed

73
Q

Factors affecting Rate

A
  • nature of the reactant/bonds
  • surface area exposed
  • concentrations
  • temperature
  • catalyst
74
Q

Law of Mass Action

A

the rate of the chemical reaction is proportional to the product of the concentrations of the reactant (k is the specific rate constant that is only affected by temp)

75
Q

Large Keq

A

equilibrium does not occur until the conc of the original reactants are small and those of the products are large

76
Q

Small Keq

A

equilibrium occurs almost as once and and relatively little product is produced

77
Q

Acid disassociation constant

A

Ka; incorporates the conc of water

78
Q

Small Ka

A

conc of un-ionized molecule must be relatively large compared with ions conc

79
Q

Large Ka

A

conc of ions are relatively high

80
Q

pH

A
81
Q

% dissociation

A

moles or liters that disassociate/original conc *100

82
Q

Ksp

A

incorporates conc of solute

83
Q

Second Law of Thermodynamics

A

entropy of the universe increases for any spontaneous process

84
Q

Gibbs free-energy equation

A

∆G = ∆H -T∆S (where temp is in K) NOT C°

85
Q

Characteristics of Common Acids

A

Water/aq solutions of acids conduct electricity

2) Acids will react with metals that are more active than hydrogen ions to liberate hydrogen
3) Acids have the abilit to change the color of indicators (such as Litmus)
4) Acids react with bases so that the properties of both are lost to form water and a salt
5) Acids react with carbonates to release carob dioxide

86
Q

Arrhenius Theory

A
  • Acids: a substance that yields hydrogen ions in an aq solution
  • Bases: a substance that yields hydroxide ions in an aq solution
87
Q

Characteristics of Common Bases

A

Bases are conductors of electricity in aq solutions
Bases cause a color change in indicators
Bases react with acids to neutralize each other and form a salt and water
Bases react with fats to form a class of compounds called soaps
Aq Solutions of bases feel slippery, and the stronger bases are very caustic to the skin

88
Q

The Bronsted-Lowry Theory

A
  • Acids: proton donors

- Bases: proton acceptors

89
Q

Conjugate acids and bases

A
  • acids give away a proton to become a conjugate base
  • bases accept a proton to become a conjugate acid
  • the stronger the base, the weaker the conjugate acid
  • the stronger the acid. the weaker the conjugate base
90
Q

Lewis Theory

A
  • Acid: electron-pair acceptor

- Bases: electro-pair donor

91
Q

Titration

A

M acid × V acid = M base × V base

92
Q

5 Methods for Preparing Salts

A 
Neutralization
Single replacement
Direct combo of elements
Double replacements
Reaction of a metallic oxide with a nonmetallic oxide
93
Q

amphoteric substances

A

substances that donate protons in the presence of strong bases and accept protons in the presence of strong acids

94
Q

Acid Rain

A

forms when oxygen, sulfur, and nitrogen combine in the atmosphere

95
Q

Buffer solutions

A

resists changes in pH

96
Q

Electrolyte

A

a substance that dissolves in water to form a solution that will conduct an electric current (if they don’t conduct…non electrolytes)

97
Q

ionization

A

ions formed by breaking the covalent bonds

98
Q

dissociation

A

ionic lattices come apart

99
Q

Cations and Anions

A
  • Cations: + (smaller than neutral atom)

- Anions: - (larger than neutral atom)

100
Q

Tip: “OILRIG”

A

Oxidation Is Lost… Reduction Is Gained

-Dr. Brennan

101
Q

Oxidation-Reduction

A

a substance that is oxidized during the reaction is separated from the substance that is reduced during the reaction; transfer of electrical energy NOT heat

102
Q

“An Ox and a Red Cat”

A

anode is oxidation and reduction is cathode

103
Q

electrochemical cell

A

system of electrodes and electrolytes in which either chemical reactions produce electrical energy or an electric current produced chemical change

104
Q

voltaic/galvanic cell

A

redox reaction occurs naturally

105
Q

allotropic

A

the existence of 2 or more different physical forms of a chemical element

106
Q

Sulfuric Acid

A

acid salts; normal salts; ionization is more extensive; hydrogen sulfide is colorless other sulfides have colors

107
Q

Halogen Family

A

colors at room temp; acceptors of electrons; Fluorine is most active (see pg 300 Barrons)

108
Q

Metals

A

luster, conduct heat and electricity, malleable, ductile, silvery color most times, not soluble without chemical change

109
Q

Carbon Dioxide

A
  • in lab usually react carbon carbonate with hydrochloric acid
  • test by passing it through limewater; cloudy if present
  • stalagmites and stalactites form
  • important for photosynthesis
110
Q

Alkane Series

A

CnHn+₂ (hydrocarbon series) methane, ethane, propane, n-Butane…etc

111
Q

alkyl halides

A

alkanes can be reacted with halogens so that hydrogens are replaced by a halogen atom

112
Q

Alkene Series

A

-double covalent bond between 2 adjacent carbon atoms; CnH₂n

113
Q

Alkyne Series

A

CnH₂n-₂ triple covalent bond between two adjacent carbons

114
Q

Aromatics

A

unsaturated ring structures; CnH₂n-₆

115
Q

Isomers

A

same formula, different structures

116
Q

Alcohols (Methanol and Ethanol)

A

functional group; methanol is simplest alcohol; ethanol is best known and most used alcohol

117
Q

Aldehydes

A

carbon double bonded to oxygen and single bonded to hydrogen and attached to an alkyl group

118
Q

Organic acids

A

Carbon double bonded to oxygen and single bonded to hydroxide (Generally R-COOH)

119
Q

Ketones

A

R-C-R and C is double bonded to Oxygen

120
Q

Ethers

A

dehydrated with sulfuric acid R-O-R

121
Q

Amines and Amino Acids

A

NH₂⁻ found in amide ions and the amino group; contain 1 or more amino groups; amide functional group is -N single bonded to 2 H’s

122
Q

Esters

A

compared to inorganic salts; the functional group is R-C-O-R where C is double bonded to another Oxygen

123
Q

Common Lab Tools

A

gravimetric balance with direct readings; pH meters; Spectrophotometer; Computer-assisted labs

124
Q

Qualitative Tests of Some Metals

A

flame tests; hydrogen sulfide tests

125
Q

Boiling point

A

increased by .51°C for every mole

SAT II: Chemistry (1 decks)

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