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General Chemistry 1 > Final Exam Practice > Flashcards

Final Exam Practice Flashcards

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1
Q

Mixture

A

A material that can be separated by physical means into two or more substances.

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2
Q

Heterogeneous Mixture

A

A mixture that consists of physically distinct parts, each with different properties.

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3
Q

Homogeneous Mixture

A

As called a solution. A mixture that is uniform in its properties throughout given samples.

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4
Q

Phase

A

One of several different homogeneous materials present in the portion of matter under study.

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5
Q

Period

A

The elements in any one horizontal row of the periodic table.

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6
Q

Group

A

The elements in any one column of the periodic table.

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7
Q

Chemical Equation

A

The symbolic representation of a chemical reaction in terms of chemical formulas.

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8
Q

Reactant

A

A starting substance in a chemical reaction.

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9
Q

Product

A

A substance that results from a reaction.

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10
Q

Mole (mol)

A

The quantity of a given substance that contains as many molecules or formula units as the number of atoms in exactly 12 g of carbon-12.

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11
Q

Avogadro’s Number

A

The number of atoms in a 12 g sample of carbon-12.

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12
Q

Molar Mass

A

The mass of one mole of the substance.

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13
Q

Percentage Composition

A

mass % A = (mass of A / mass of whole) x 100%

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14
Q

Stiochiometry

A

The calculation of the quantities of reactants and products involved in a chemical reaction.

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15
Q

Limiting Reactant

A

The reactant that is entirely consumed when a reaction goes to completion.

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16
Q

Theoretical Yield

A

The maximum amount of product that can be obtained by a reaction from given amounts of reactants.

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17
Q

Percentage Yield

A

% yield = (actual yield in experiment / theoretical yield in calculations) x 100%

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18
Q

Molecular Equation

A

A chemical equation in which the reactants and products are written as if they were molecular substances, even though they may actually exist in solution as ions.

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19
Q

Complete Ionic Equation

A

A chemical equation in which strong electrolytes are written as separate ions in solution.

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20
Q

Spectator Ion

A

An ion in an ionic equation that does not take part in the reaction (it is cancelled out).

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21
Q

Net Ionic Equation

A

An ionic equation from which spectator ions have been cancelled.

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22
Q

Oxidation Number

A

The actual charge of the atom if it exists as a monatomic ion, or a hypothetical charge assigned to the atom in the substance by simple rules.

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23
Q

Oxidation-Reduction Reactions

A

A reaction in which electrons are transferred between species or in which atoms change oxidation numbers.

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24
Q

Half-Reaction

A

One of two parts of an oxidation-reduction reaction. One part includes the loss of electrons, the other gains electrons.

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25
Q

Oxidation

A

The loss of electrons leads to an increase in oxidation number.

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26
Q

Reduction

A

The gain of electrons leads to a decrease in oxidation number.

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27
Q

The Scientific Method

A
  1. State the problem
  2. Hypothesis (tentative statement)
  3. Experiments (observations carries out in a controlled manner.
  4. Collect and analyze data (determine relationships of data)
  5. Draw Conclusions
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28
Q

Mega (M)

A

1x10^6

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29
Q

Kilo (k)

A

1x10^3

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30
Q

Deca (d)

A

1x10^-1

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31
Q

Centi (c)

A

1x10^-2

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32
Q

Milli (m)

A

1x10^-3

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33
Q

Micro (µ)

A

1x10^-6

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34
Q

Nano (n)

A

1x10^-9

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35
Q

Pico (p)

A

1x10^-12

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36
Q

Atomic Mass Unit (amu)

A

An amu is equal to half the mass of isotope carbon-12

1 amu = 1.6606x10^-24 g

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37
Q

SI Units

A
  1. length: meter, m
  2. mass: kilogram, kg
  3. time: second, s
  4. temperature: Kelvin, K
  5. amount of a substance: mole, mol
  6. electric current: ampere, A
  7. luminous intensity: candela, cd
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38
Q

Derived Units

A
  1. area: m^2
  2. volume: m^3
  3. density: kg / m^3
  4. speed: m/s
  5. acceleration: m / s^2
  6. force: kg x m / s^2 (N)
  7. pressure: kg / m x s^2 (Pa)
  8. energy: kg x m^2 / s^2 (J)
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39
Q

Mass Number

A

Total number of protons and neutrons in the nucleus.

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40
Q

Oxidation States

A
  1. The oxidation number of an atom in elemental form is 0.
  2. The oxidation number of a monatomic ion is equal to its charge.
  3. The oxidation number of fluorine (F) is always -1.
  4. The oxidation number of Group One atoms is +1.
  5. The oxidation number of Group Two atoms is +2.
  6. The oxidation number of Group Seventeen atoms is -1 wen combined with elements to the left or below it on the periodic table.
  7. The oxidation number of oxygen (O) is usually -2, except with fluorine (+), peroxides (-1), and superoxides (-.5).
  8. The oxidation number of hydrogen (H) is -1 with metals, but +1 with nonmetals.
  9. The sum of the oxidation number for all atoms in a compound must equal 0.
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41
Q

Atomic Mass

A

The average atomic mass for the naturally occurring element, in amu.

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42
Q

Periodic Table

A

Created by Mendeleev and Meyer, it is a tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements. It is arranged by atomic number.

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43
Q

Group 1A

A

Alkali Metals

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44
Q

Group 7B

A

Halogens

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45
Q

Group 8

A

Noble Gases

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46
Q

Lanthanides

A

The first of the lower rows of inner transition metals.

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47
Q

Actinides

A

The second of the lower rows of inner transition metals.

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48
Q

Group 2B

A

Alkaline Earth Metals

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49
Q

Group 6B

A

Chalcogens

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50
Q

-ate

A

This tells us that there is a greater number of oxygen ions in the compound. The acid suffix is -ic.

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51
Q

-ite

A

This tells us that there is a lesser amount of oxygen ions in the compound. The acid suffix is -ous.

52
Q

Acid Nomenclature

A
  1. When there is no oxygen in the acid: hydro- + root name + -ic + acid
  2. When the acid contains oxygen: root name + -ic + acid
  3. When there are only two oxygen containing species (smaller oxygen content): root name + -ous + acid
  4. When there is a higher oxygen content: root name + -ic + acid
  5. When there is less oxygen than in the -ous acid: hypo- + root name + -ous + acid
  6. When there is more oxygen than in the -ic acid: per- + root name + -ic + acid
53
Q

Base Nomenclature

A
  1. When the base contains OH- ions: metal name + hydroxide

2. When there is no OH-: there is a common name, i.e. ammonia, methylamine.

54
Q

Covalent Compounds Nomenclature

A
  1. For the first element: add a prefix unless there is only one (di, tri, tetra, etc)
  2. For the second element: add a prefix related to how many there are (mono, di, tri, tetra, etc)
55
Q

Hydrate Nomenclature

A

Salt name + prefix + hydrate

56
Q

Balancing Net Ionic Equations

A
  1. Balance equation.
  2. Break into ions if soluble (see solubility rules).
  3. Cancel any spectator ions.
  4. Rewrite with non-spectator ions.
  5. Recheck

If everything cancels, there is no reaction.

57
Q

Balancing Redox Reactions

A

For each half reaction:

  1. Balance all non-oxygen and non-hydrogen species first.
  2. Balance oxygen and hydrogen by adding H+ and/or H2O if acidic or OH- and/or H2O if basic.
  3. Balance charge by adding electrons.
  4. Add half reactions together and cancel common species.
  5. Recheck.

Acidic: Add H+ to side with higher oxidation state
Basic: Add OH- to side with lower oxidation state

58
Q

Titration

A

The process of determining the concentration of one substance whose concentration is unknown by reacting it with a solution of another substance whose concentration is known.

Molarity of known -> moles of known -> moles of unknown -> molarity of unknown

59
Q

Dilution

A

M1V1 = M2V2

60
Q

Molarity (M)

A

Moles of a substance / Liters of a substance

61
Q

Molality (m)

A

Moles of solute / kg of solvent

62
Q

q

A

This is the symbol for heat. Defined as the transfer of thermal energy between two substances which are at different temperatures.

q = ms∆t
q = C∆t

A positive q value represents an endothermic process. A negative q value represents an exothermic process.

63
Q

∆H

A 
∆H = enthalpy 
∆H(rxn) = H(products) - H(reactants)
∆H(rxn) = ∑∆H(products) - ∑∆H(reactants)

When ∆H is positive, heat is added to the system (endothermic). When ∆H is negative, heat is removed from the system (exothermic).

64
Q

Exothermic Process

A

Process that gives off heat from system to the surroundings.

T increases
∆E is negative
q is negative
∆H is negative

65
Q

Endothermic Process

A

Heat is supplied (added) to the system.

T decreases
∆E is positive
q is positive
∆H is positive

66
Q

Thermochemistry Equations/Conversions

A

∆E = E(final) - E(initial)
q = ms∆t
heat capacity: C = mS

1 kilocalorie = 1 calorie
1 calorie = 4.184 J
1 Nüt cal = 1 kcal

67
Q

Photoelectric Effect

A

Procedure completed in 1905 by Einstein. He concluded that light is both a particle and a wave.

68
Q

Discharge Tube Experiments

A
  1. Construction by Humprey Daviy in 1821. The negative electrode is called a cathode and the positive electrode is called an anode. Also had a vacuum arm at the bottom.
  2. Discoveries: Electricity flow from one electrode to other, different gases produce different colors, and no gas gives off green light because of cathode rays.
69
Q

Variation of Tube Experiment

A

A zinc sulfide plate was inserted with the shadow on the cathode side. Shows rays coming from cathode are negatively charged.

70
Q

Plum Pudding Model

A

JJ Thompson.
Two cathode leaves, charged with cathode rays. The cathode leaves repel each other and the temperature increases.
Electroscope magnetic fields: Cathode repels negative plate and attracted to positive side.
Cathode with holes: cathode is attracted to anode
Conclusion: Canal rays are heavier and have a positive charge.

71
Q

Metal Foil Experiment

A

EJ Rutherford
Observations: most particles went through the foil, others scattered.
Conclusion: atom is mostly void space, positive charge concentration is in a tiny volume with large mass relative to surroundings.

72
Q

Theory of Hydrogen Atoms

A

Niels Bohr in 1913.

  1. Electrons around nucleus go in a circular orbit.
  2. No energy is lost or gained if the electron stays in orbit.
  3. Energy changes if the electron changes orbit.
  4. Orbital angular momentum of electron is quantized.
73
Q

Quantum Equations

A 
c = (wavelength) x (frequency)
E = h x frequency 
∆E = RH x ((1/n^2) - (1/n^2))
74
Q

Nature of Light

A

James Maxwell in 1873 said that light consists of waves.
Einstein in 1905 said light consists of waves and particles.
The speed of light is 2.998x10^8 m/s (c)

75
Q

Wave Equation and probalility

A

E. Schrödinger in 1926 said the wave equation is the most probably region to find electrons in (atomic orbital)

Psi is the wave function (n, l, m, s)
Psi squared is the probability of finding electrons some distance (r) away from the nucleus.

76
Q

Quantum Number: n

A

Principle quantum number
Value: 1, 2, 3, … ∞
Physical significance: energy or size of electron orbital

77
Q

Quantum Number: l

A

Angular or azimuthal
Value: 0 to (n-1)
Physical Significance: shape or energy of orbital

78
Q

Quantum Number: m

A

Magnetic
Value: -l to +l
Physical Significance: orientation

79
Q

Quantum Number: s

A

Spin
Value: +.5, -.5
Physical Significance: none

80
Q

Ionization Energy

A

Generally decreases down the group.

Generally increase across the period.

81
Q

Atomic Size

A

Generally increase down the group.
Generally decrease across the period.

Different from all other periodic trends.

82
Q

Electronegativity

A

Generally decreases down the group.

Generally increases across the period.

83
Q

Electron Affinity

A

Generally decreases down the group.

Generally increases across the period.

84
Q

Predicting Reaction Products

A
  1. Identify each element in one reactant metal or nonmetal.
  2. a) Elemental metals will be oxidized by elemental nonmetal and reduce nonmetal.
    b) Elemental nonmetals will will oxidize a nonmetal lying to the left or below it in the periodic table.
    c) Metals of Group 1, Ca, Sr, and Ba can reduce water and acids to make H2 and metal.
  3. Acid-Base reaction: salt and H2O as product
  4. Metathesis: swap metals and nonmetals
85
Q

Lewis Dot Diagram Information

A

Lewis observed a) molecules have an even number of electrons and b) molecules have no unpaired electrons except for NO, NO2, and O2

Lewis Dot symbols: elemental symbol with dots related to the number of valence electrons

Octet Rule: each atom in the molecule tends to achieve eight valence electrons (2 for H and He)

One pair of electrons is a single bond.
Two pairs of electrons is a double bond.
Three pairs of electrons is a triple bond.

C, N, O, P, and S usually form multiple bonds.

86
Q

Writing Lewis Dot Diagrams

A
  1. Write a skeletal structure of compound showing which atoms bond to each atoms.
  2. The least electronegative element is the center atom.
  3. Count the total number of valence electrons present.
    a) For polyatomic anions, add the number of electrons of negative charge to total.
    b) For polyatomic cations, subtract the number of electrons of positive charge from total.
87
Q

Resonance Structures

A

Double or triple bond changes places. This shoes that the bond is moving.

88
Q

Formal Charge

A

(total number of valence electrons in the atom) - (total non-bonding electrons) - .5(total bonding electrons)

89
Q

Exceptions to Octet rule

A
  1. Less than Octet: Group 2, 3
  2. Odd electron molecules: NO, NO2
  3. More than Octet: have 3d
90
Q

Bond Energy

A

The energy needed to break chemical bonds in gaseous state.

∆H = ∑D(reactants) - ∑D(products)

91
Q

Bond Order

A

.5(bonding electrons - anti-bonding electrons)

92
Q

Paramagnetic

A

A substance that is weakly attracted by a magnetic field because of unpaired electrons.

93
Q

Diamagnetic

A

A substance that is not attracted by a magnetic field because it has no unpaired electrons.

94
Q

Linear

A

AX2, 180º

95
Q

Trigonal Planar

A

AX3, 120º

96
Q

Bent

A

AX2E, less than 120º

97
Q

Tetrahedral

A

AX4, 109.5º

98
Q

Trigonal Pyramidal

A

AX3E, less than 109.5º

99
Q

Bent

A

AX2E2, less than 109.5º

100
Q

Trigonal Bipyramidal

A

AX5, 90º and 120º

101
Q

Teetor Totter

A

AX4E, less than 90º and less than 120º

102
Q

T-Shaped

A

AX3E2, less than 90º

103
Q

Linear

A

AX2E3, 180º

104
Q

Octahedral

A

AX6, 90º

105
Q

Square Pyramidal

A

AX5E, less than 90º

106
Q

Square Planar

A

AX4E2, 90º

107
Q

The Hybridization Concept

A
  1. Not applied to isolated atoms; it is used to explain bonding in a molecule.
  2. Hybrid orbitals have different shapes than atomic orbitals.
  3. It requires an input of energy.
  4. Covalent bonds in polyatomic molecules form by the overlap of hybridized orbitals with other hybridized orbitals or the overlap of hybridized orbitals with unhybridized orbitals.
108
Q

Hybridization

A
  1. AX2: sp
  2. AX3: sp2
  3. AX4: sp3
  4. AX5: sp3d
  5. AX6: sp3d2
109
Q

Boyle’s Law

A

P1V1 = P2V2

110
Q

Charles’ Law

A

(V1/T1) = (V2/T2)

111
Q

Avogadro’s Law

A

V is proportional to n

112
Q

Amonton’s Law

A

P is proportional to T

113
Q

Combined Gas Law

A

(P1V1 / T1) = (P2V2 / T2)

114
Q

Ideal Gas Law

A

PV = nRT

115
Q

Graham’s Law

A

(rate A / rate B) = (time B / time A) = (√(MM of B) / √(MM of A))

116
Q

Dalton’s Law

A

P(T) = P(gas A) + P(gas B) …

X = mole ratio = (partial P / total P)

117
Q

Van der waals Equation

A

(P + (n^2a / V^2)) x (V - nb) = nRT

118
Q

Kinetic Molecular Theory of Gases

A
  1. Gas molecules move in random motion.
  2. Molecules hit the walls of their container, creating pressure.
  3. Gas molecules collide elastically (there is no loss in energy)
  4. At a low pressure, the distance between the molecules is large compared to the size of the molecules.
  5. At a low pressure, attraction can be ignored.
  6. Molecules can be considered to have no volume because of #4.
  7. The average kinetic energy of all the gas molecules at the same temperature is the same.
119
Q

Nuclear Reactions

A
  1. alpha particle emission
  2. Beta particle emission
  3. positron emission
  4. electron capture
  5. gamma emission
120
Q

Half-Life

A

The time required for half the atoms in a sample to decay.

k = 0.693 / t(half life)
k = (1/t) x ln(No/Nt)
121
Q

Nuclear Binding Energy

A

Nucleons are protons and neutrons.

E = mc^2

122
Q

Soluble Compounds

A
  1. Group 1
  2. NH4+
  3. Acetates
  4. Nitrates
  5. Chlorides (except Ag, Hg, Pb)
  6. Bromides (except Ag, Hg, Pb)
  7. Iodides (except Ag, Hg, Pb)
  8. Sulfates (except Ca, Sr, Ba, Ag, Hg, Pb)
  9. Soluble salts
  10. Strong acids (HCl, HBr, HI, HNO3, HClO4, H2SO4, H3PO4)
  11. Strong bases (NaOH, KOH, Ba(OH)2)
123
Q

Insoluble Compounds

A
  1. Carbonates (except Group 1, NH4+, Ca, Sr, Ba)
  2. Phosphates (except Group 1, NH4+, Ca, Sr, Ba)
  3. Sulfides (except Group 1, NH4+, Ca, Sr, Ba)
  4. Hydroxides (except Group 1, NH4+, Ca, Sr, Ba)
  5. Weak electrolytes
  6. Covalent compounds H2O, NH3)
  7. Weak acids (HF, H2CO3, H2SO3, HCN, HC2H3O2, H2S)
  8. Gases at room temperature
124
Q

Radioactive Dating

A

Half life of carbon-14 is 5730 years.

125
Q

Nuclear Fission

A

Heavy nucleus splits (spontaneously or upon contact) with the release of energy.

126
Q

Nuclear Fusion

A

Atomic nuclei of low atomic number fuse to form heavier nucleus, with the release of energy.

127
Q

Nuclear Stability

A
  1. Odd protons and odd neutrons is least stable.
  2. Odd protons and even neutrons is more stable.
  3. Even protons and odd neutrons is even more stable.
  4. Even protons and even neutrons is most stable.

General Chemistry 1 (1 decks)

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