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Chemistry > Semester 1 Exam > Flashcards

Semester 1 Exam Flashcards

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1
Q

List metric prefixes pico-giga

A 
pico p 10^-12
nano n 10^-9
micro μ 10^-6
milli m 10^-3
centi c 10^-2
kilo k 10^3
mega M 10^6
giga G 10^9
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2
Q

Temperature conversions

A

C=((5/9)(F-32)
F=((9/5)C)+32
K=C+273

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3
Q

How to make an isotope symbol

A

Mass number up left
Element symbol
Atomic number down left

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4
Q

Acid naming w/ polyatomics

A

I -ate something -ic(ky) but I didn’t (b)-ite a (m)-ous
eg carbonate=carbonic acid
nitrate=nitrous acid

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5
Q

Acid naming w/ binary acids (ie H+ion)

A

Hydro+ion name+ic

eg HCl=hydrochloric acid

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6
Q

What is an alkane

A

A hydrocarbon with all single bonds

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7
Q

How to make alkane formulas

A 
C(n)H(2n+2)
1 C=Methane
2 C=Ethane
3 C=Propane
4 C=Butane
5+ C=Greek prefixes (eg Pentane
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8
Q

How to demonstrate heat added to a chemical equation

A

Put a Δ over the arrow

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9
Q

Strong electrolytes

A

Strong acids: HCl HBr HI HNO3 HCLO3 HCLO4 H2SO4
Strong bases: Group 1 hydroxides
Ionic compounds

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10
Q

Weak electrolytes

A

weak acids and bases

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11
Q

NH3 + H2O–>

A

NH4 + OH-

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12
Q

Definition of acid

A

Molecule that can donate a proton

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13
Q

Base

A

Molecule that accepts protons and creates hydroxides

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14
Q

What is reduction

A

Gain e- reduce oxidation #

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15
Q

What is oxidation

A

Lose e- increase oxidation #

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16
Q

Oxidation rules

A 
Element by itself: 0
Neutral compounds add up to 0
Ions add up to charge of ion
Oxygen: 2-
Hydrogen 1+
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17
Q

Kinetic Energy=

A

1/2(mass x velocity)

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18
Q

Joules calorie

A

4.184J

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19
Q

1st law of thermodynamics

A

Conservation of energy

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20
Q

How is enthalpy notated

A

ΔH

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21
Q

How is enthalpy calculated

A

per mole

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22
Q

What does the sign of enthalpy indicate

A

Positive:endothermic
Negative:exothermic

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23
Q

q=CΔT purpose and parts

A

Heat of system when taking into account calorimeter, C:heat capacity of calorimeter

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24
Q

Heat capacity

A

Specific heat w/o mass

Unit J/ºC

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25
Q

Enthalpy of formation

A

E to form a compound, generally negative because formation is exothermic

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26
Q

How to calculate ΔH of reaction

A

(ΔHf of products) - (ΔHf of reactants)

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27
Q

If you add heat to ice at 0ºC what happens to temperature

A

It stays the same;the energy is used to melt the ice

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28
Q

How to calculate energy to melt and vaporize

A

ΔHfusion x mass = energy to melt

ΔHvaporization x mass=energy to vaporize

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29
Q

Formula for speed of light

A

Wavelength x Frequency

λ v = c

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30
Q

Wavelength of visible light

A

400-700nm

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31
Q

List types of electromagnetic radiation in order of increasing wavelength

A

gamma, x-rays, UV, visible, infrared, microwaves, radio

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32
Q

E of photon =

A

Planck’s constant x frequency

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33
Q

Heisenberg’s uncertainty principle

A

You can’t simultaneously know the velocity and position of a particle

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34
Q

Quantum number

n

A

n: principle quantum #

~Integers ≥1, usually [1, 4]

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35
Q

Quantum number

l

A 
l: angular momentum quantum #
~[0-(n-1)]
~l= 0 letter s
l=1 letter p
l=2 letter d
l=3 letter f
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36
Q

Quantum number

ml

A

ml: magnetic quantum #
~[-L, +L]
Determines position in orbital

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37
Q

Quantum number

ms

A

ms: spin magnetic quantum #
~(-1/2) or (+1/2)
- = down arrow + = up arrow

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38
Q

Shape of s orbital

A

Sphere

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39
Q

Shape of p orbital

A

dumbbell

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40
Q

Shape of d orbital

A

Clover or dude with hula hoop

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41
Q

Exceptions when doing electron configurations

A

ns2 nd4 → ns1 nd5

ns2 nd9 → ns1 nd10

42
Q

What is effective nuclear charge

A

Zeff= positive charge of nucleus after taking into account electron screening

43
Q

How to calculate Zeff

A

Zeff= Protons (Z) - Core Electrons (S)

44
Q

What is electron affinity

A

Energy change to add an electron to a gaseous atom

Almost always negative

45
Q

What is ionization energy

A

Energy change to remove an electron from a gaseous atom

Almost always positive

46
Q

What is an isoelectronic series

A

Group of ions all containing the same number of electrons
Higher atomic #=lower atomic radius
Higher atomic #=higher Zeff

47
Q

Trend for metallic character

A

Increases left and down PT

48
Q

Metal oxide + water –>

A

Metal hydroxide (base)

49
Q

Nonmetal oxide + water–>

A

acid

50
Q

Nonmetal oxide + base–>

A

salt + water

51
Q

alkali metal + water–>

A

H2 + base (explosive)

52
Q

What group is alkali metals

A

1

53
Q

What group is alkaline earth metals

A

2

54
Q

What do lewis symbols show

A

Just the one atom and its valence e-

55
Q

When are bonds polar

A

When one atom is more electronegative than the other

56
Q

Trend of electronegativity

A

Increases up and right on PT

57
Q

Which way does the dipole moment arrow point

A

Towards the more electronegative atom in the bond

58
Q

Exceptions to the octet: less than the octet

A

B=6 e-
Be=4 e-
H= 2 e-

59
Q

Exceptions to the octet: more than the octet

A

Below 1st two periods

60
Q

How to find formal charge

A

Draw lewis structure
(total ve-) - (# bonds + # unbonded e-)
nb double bonds count as 2
Atoms w/ higher electronegativity more likely to take a negative charge, oxygen often willing to take 1- charge

61
Q

What is an isomer

A

Same composition, different structure

62
Q

What is a resonance structure

A

Same composition and structure, different bonds

63
Q

What does VSEPR stand for

A

Valence shell electron pair repulsion theory

64
Q

Steps for predicting shape w/ VSEPR

A

~Draw the lewis structure and count the number of electron domains around the central atom
~Determine the electron domain geometry by arranging the electron domains around the central atom so that the repulsions among them are minimized
~Use the arrangement of the bonded atoms to determine the molecular geometry

65
Q

Hybrid orbitals: 1s + 3p=

A

4sp3

66
Q

What types of bonds are sigma vs pi

A

All single bonds are sigma
Double bonds have one sigma one pi
Triple bonds have one sigma two pi

67
Q

Pressure=

A

force/area

68
Q

Pressure units and conversions

A

1 atm=101325 Pa
1 atm=101.325 kPa
1 atm=760 torr=760mmHg

69
Q

Different values of R in PV=nRT

A

R=.0821 for atm

R=62.4 for torr

70
Q

Gas Laws: Boyle’s law

A

P1V1=P2V2

71
Q

Gas Laws: Charles’ law

A

V1/T1 = V2/T2

72
Q

Gas Laws: Avogadro’s law

A

V1/n1 = V2/n2

73
Q

Gas Laws: Ideal gas law

A

PV=nRT

74
Q

How to represent direct proportionality

A

X / Y

75
Q

How to represent inverse proportionality

A

XY

76
Q

Density=

A

mass/volume
d=PM / RT
M=molar mass

77
Q

STP

A

1 atm 0ºC

78
Q

1 mol gas @ STP

A

22.4L

79
Q

Parts of Kinetic molecular theory

A

~Gases consist of large numbers of molecules that are in continuous random motion
~The combined volume of all the molecules of the gas is negligible compared to the volume of its container
~IMFs between gas molecules are negligible
~The average kinetic energy of the molecules in a gas does not change w/ time
~Average kinetic energy is proportional to absolute temperature

80
Q

What does it mean that average kinetic energy is proportional to absolute temperature

A

At any given temperature the molecules of all gases have the same average kinetic energy

81
Q

What is effusion

A

Rate at which a gas can pass through a pinhole

82
Q

How to calculate rate of effusion

A

r1/r2 = √M2 / M1
r=rate
M=molar mass

83
Q

What causes gases to be less ideal

A

Gases are less ideal as mass increases and temperature decreases

84
Q

What condition needs to be in place for hydrogen bonding

A

Hydrogen must be bonded to N O or F

85
Q

What is capillary action

A

Capillary action causes water to rise up the walls of a small diameter glass tube

86
Q

How do adhesive and cohesive forces affect capillary action

A

~Adhesive forces between the liquid and the wall tend to increase the surface area of the liquid
~The cohesive force of surface tension tends to reduce surface area, thereby pulling the liquid up the tube

87
Q

What stops capillary action

A

The liquid climbs until the force of gravity on the liquid balances the adhesive and cohesive forces

88
Q

What are the triple and critical points

A

On a phase change diagram:
~The triple point is the point where the lines separating solid, liquid, and gas phase meet
~The critical point is the point after which the substance becomes a supercritical fluid

89
Q

What are the conditions for normal boiling point

A

1 atm

90
Q

What is the most general rule of solubility

A

Like dissolves like–ie polar dissolves polar and nonpolar dissolves nonpolar

91
Q

How to supersaturate a solution

A

Add heat to increase the solubility of the solute

92
Q

How do IMFs affect solubility

A

Strong IMF=low solubility

93
Q

How does pressure affect gas solubility

A

High pressure=high gas solubility

94
Q

How does temperature affect liquid and gas solubility

A

High temp=high liquid solubility

High temp=low gas solubility

95
Q

Units of concentration: Molarity

A

mols solute / L solution

96
Q

Units of concentration: molality

A

mols solute / kg solvent

97
Q

Units of concentration: Mass %

A

g solute / g solution

98
Q

What do colligative properties depend on

A

Concentration

99
Q

What happens to the boiling and freezing points of a solvent when a solute is added

A

Boiling point elevates

Freezing point depresses

100
Q

Formula for boiling pt elevation and freezing point depression

A

ΔT = i m k
ΔT: add to boiling, subtract from freezing
i=VanHoff factor, number of pieces substance breaks into when dissolved
m=molality
k=boiling/freezing point constant of solvent

101
Q

On what does a state function depend

A

The state of the material (ie solid, liquid or gas)

102
Q

What is the effect of long hydrocarbon chains/nonpolar tails

A

Solubility in water decreases

Chemistry (4 decks)

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