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Malika Gen Chem MCAT > Gen Chem Terms > Flashcards

Gen Chem Terms Flashcards

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1
Q

A balanced reaction has the same number of _______ of each type on both sides and the same net ______ on both sides.

A

atoms, charge

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2
Q

Is there an error in the balancing of this reaction?

A

Yes, not the same # of H atoms on each side, and not the same net charge.

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3
Q

Calculating theoretical yield question

A
  1. Balance equation.
  2. Convert each starting reactant to moles of calcium citrate to find the limiting reactant
  3. Convert the smaller amt. of moles to grams to get the MAXIMUM product of calcium citrate produced.
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4
Q

Calculating percent yield

A

THEORETICAL YIELD - The answer you get from determining the limiting reactant
Actual yield - usually given in the question

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5
Q

Bronsted Lowry acid vs Lewis Acid definitions

A

Bronsted Lowry acid - molecule which donates a proton (H)
Lewis acid - molecules which accepts electrons

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6
Q

Question about BL acid and Lewis acid

A

H2NO3 acts as an H+ donor, so it is considered a Bronsted-Lowry Acid.
However, for the reverse reaction, NO2 is an electron pair acceptor because it acts as an electrophile ONLY.

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7
Q

In a coordinate covalent bond, the central atom (typically a metal), acts as a Lewis ________, while the connecting atoms act as Lewis ________.

A

acid, base
Central metal ion accepts electron lone pairs, while the connecting atoms donate their electron lone pairs

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8
Q

Difference between noncovalent IMFs

A

Ion-dipole: a fully charged ion is attracted to an atom with a partial charge
Dipole-induced dipole: attraction between a polar molecule & nonpolar molecule, which induces a temporary dipole
Dipole-dipole: attraction between atoms of opposite partial charges
London dispersion: nonpolar bonds induce temporary weak dipoles

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9
Q

Dipole-induced dipole IMF

A
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10
Q

Properties of metals

A

Small ionic radius, low electronegativity, low ionization energy, low electron affinity, great electrolytes

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11
Q

Examples of coordinate covalent bonds

A

-Multiple NH3 molecules bonded to a transition metal:
Cu(NH3)4^2+, Co(NH3)4^2+,
NH3BF3
-Al2Cl6
-Ni(PBr3)4^2+

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12
Q

What defines a molecule to be the most polar? Out of all the halogens which one forms the most polar bond w/ H?

A

-Highest electronegativity
-Smallest atomic radius
-High ionization energy
Fluorine forms the most polar bond with H.

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13
Q

Metallic character increases as you go ______ the periodic table.

A

down

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14
Q

Condosity

A

The molar concentration of NaCl that has the same specific electric conductance as the solution

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15
Q

Ionic compound exceptions

A

NH4Cl, (NH4)2SO4, NH4BF4

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16
Q

Ranking compounds based on ionic character

A

-The higher difference between electronegativities between 2 atoms, the more ionic it is
-More metallic metal (further down in table) in the compound, makes it more ionic

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17
Q

Characteristics of covalent bonds

A

-Have electronegativity difference between 0.5-1.7
-Electrons are shared equally or unequally depending on type of covalent bond
-Involved 2 NONMETALS

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18
Q

For a compound to be considered ionic, the electronegativity difference should be _____________.

A

greater than 1.7

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19
Q

What is meant by the most stable electron configuration?

A

The electron configuration of the atom in its ground state.

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20
Q

gamma radioactive decay

A

-Caused by release of a high energy photon
-Present in EVERY type of decay
-DOESN’T alter atomic mass or atomic number

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21
Q

alpha radioactive decay

A

-Caused by release of helium nucleus
-Reduces atomic mass by 4 and atomic number by 2

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22
Q

B+ decay (positron emission)

A

Positron emission - atomic # decreases by 1; proton converts to a neutron & ejects a positron

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23
Q

Beta decay: electron capture

A

Electron capture - atomic # decreases by 1; proton captures an electron to convert it into a neutron

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24
Q

B- decay (electron emission)

A

Electron emission - atomic # increases by 1; a neutron is converted into a proton, thus emitting an electron

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25
Q

what happens molecularly during fluorescence

A

The solutions change color due to electrons absorbing energy and being promoted to a higher energy level. Once the electrons are relaxed, they return back to the original energy level, emitting fluorescence (emit longer wavelength w/ less energy than energy absorbed).

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26
Q

Unstable compounds have a high (bond dissociation energy/heat of combustion).

A

High heat of combustion because breaking bonds of an unstable compounds is easy and it releases a lot of energy. The bond dissociation is low.

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27
Q

How to determine units of rate constant (k) when given the actual rate law?

A

Use simple equation:
M^(1 - total order of reaction) *time^-1

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28
Q

In order for an electron to be ejected, it must absorb enough energy. Which equation models this concept?

A

When an e- absorbs light, it gains kinetic energy allowing it to be ejected.
KE = E - Φ
Φ is the work function
E is the energy the electron absorbs, therefore E must be greater than Φ

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29
Q

Compare insulators & conductors

A
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30
Q

Flashpoint

A

The temp at which a metal decomposes or explodes

31
Q

What are the subshells?

A

s, p, d, f

32
Q

Pauli exclusion principle

A

Each orbital within a subshell can hold a max of 2 electrons, but they must have OPPOSITE spins
No two electrons in an atom can have the same set of quantum #s

33
Q

Aufbau principle

A

Electrons must be filled from lower energy to higher energy subshells

34
Q

What are orbitals?

A

Orbitals are located within a subshell.
s - 1 orbital
p - 3 orbitals
d - 5 orbitals
f - 7 orbitals

35
Q

principle (n) quantum #

A

principle quantum number, describes the number of shells

36
Q

azimuthal (l) quantum #

A

describes the type of subshells (shape) an atom has
l = 0 –> s subshell
l = 1 –> p subshell
l = 2 –> d subshell
l = 3 –> f subshell

37
Q

How to calculate l

A

up to n-1
n = 2, so l = 0, 1
n = 3, so l = 0, 1, 2

38
Q

magnetic (ml) quantum #

A

specifies the orientations of the orbitals within a subshell, describes how many orbitals are within a subshell

39
Q

If l = 1, what are the ml quantum #s?

A

ml = -1, 0, 1
3 orbitals within p subshell

40
Q

Hund’s rule

A

Within a subshell, orbitals are filled such that there are a max # of half filled orbitals, BEFORE doubling up

41
Q

Notation for writing an element from the periodic table

A

Top number is atmoic mass (# of protons + neutrons)
bottom number is atomic # (only protons)

42
Q

If an isotope is written as such: Francium-223. What does the number mean?

A

223 would be their atomic mass (# of protons + neutrons).
Isotopes have different atomic masses.

43
Q

Elements within the same (group/period) of the periodic table have similar chemical properties.

A

group because they all have the same # of valence e-

44
Q

Electronegativity definition & trend

A

The tendency of an atom to attract electrons within a bond.
It increases as you go from left to right and decreases as you go down a group.

45
Q

Ionization energy definition & trend

A

The energy required to remove the 1st most loosely bound electron

Increases as you move to the right, decreases as you move down a group.

46
Q
A

Cs has the lowest ionization energy out of all, because the valence e- occupies a shell much farther from the nucleus (less tightly bound) so it is readily given up.
Therefore, it MOST reactive (unstable)

47
Q

Electron affinity definition & trend

A

The energy that results when adding an electron to an atom.

Electron affinity increases moving left to right across a period, and decreases moving down a group.

48
Q

Exceptions to electron affinity

A

Even though generally EA decreases as you go down a group, O and F have low EA (located at top of group) because they have small atomic radii which creates electron-electron repulsion forces.

49
Q

True or false:
Electrons are more easily accepted by nonmetals than by metals.

A

True

50
Q

Ionic radius definition & trend

A

Applies on to metal cations and metal anions.

Ionic radii decreases moving left to right and INCREASES moving down a group.
Group 1A has bigger ionic radii than group 2A. Group 5A has bigger ionic radii than 6A, and so on. Note that anions (nonmetals) overall have bigger ionic radii than cations.

51
Q

Effective nuclear charge definition and trend

A

The attraction that between valence electrons and the nucleus which creates a net positive charge that is experienced by the valence electrons.

Increases as you go across a period to the right. Stays the same down the group.

52
Q

Calculating effective nuclear charge

A

protons - # of innershell core electrons (nonvalence)

53
Q

Atomic radius definition & trend

A

The size of the neutral element itself including the shells around it.

Atomic radius decreases as you go across a period to the left. As you go down a group, it increases.

54
Q

How does effective nuclear charge affect atomic radius?

A

A higher effective nuclear charge causes greater attractions that the electrons experience, pulling the electron cloud closer to the nucleus which results in a smaller atomic radius.

55
Q

Figuring out the half life of a substance when given original concentration

A

For the 1st half life: divide the concentration by 2
Each time you go up a half life, you should divide the previous # by 2 (compounding)
*The time to get to the next half life is the SAME each time

56
Q

What is the oxidizing agent in this reaction?

A

It is O2, a common oxidizing agent which oxidizes MnO4. O2 gets reduced in the process and gains 2 electrons. This can be seen by tracking the reduction in oxidation # of O2.

57
Q

Hard question

A
58
Q

Low ionization energy corresponds to (low reactivity/high reactivity).

A

The lower the ionization energy the atom has, the more reactive it is.
Within a group, the atom furthest down is most reactive.

59
Q

True or false
Generally, the lower the ionization energy of a metal, the more reactive it is.

A

True, ionization energy and reactivity are inversely related.

60
Q

Properties of alkaline earth metals

A

-Oxidation state of +2
-Increasing reactivity with increasing atomic #
-When reacting w/ water, form a metal oxide and release H2 gas

61
Q

Properties of alkali metals

A

-Oxidation state of +1
-First group/column in table

62
Q

Relationship between atomic radius & strength of bond

A

Smaller atomic radius = stronger bond = less acidic

63
Q

Calculating mass percent of a component in a mixture/substance

A

mass percent = mass of component/total mass of mixture

64
Q

hard question

A
65
Q

What is polarizability? Polarizability (increases/decreases) as the size of an atom increases. Why? How does polarizability affect acidity?

A

The ease to which electron density in the atom can be distorted by an external electric field.

Increases because the larger the atom, the less the electrons are attracted to the nucleus (more shells) the more easy a dipole can be formed (electrons can be pulled away to form a bond).
The more polarizable the element, the more acidic it is.

66
Q

Hybridization of sp, sp2, and sp3 atoms correlates with what bond angles respectively?

A

180˚, 120˚, and 109.5˚

67
Q

Ideal gas assumptions:

A

-The law PV = nRT is obeyed
-Individual molecular volume & intermolecular forces are negligible
-At 0˚C & 1 atm, 1 mol of gas occupies 22.4L

68
Q

Deviations from ideal gas

A

-Low temperature
-High pressure
-High molecular volume & strong IMFs

69
Q

Calculating osmolarity of a solution

A

molarity of solution x n

n is the # of ions that the solution dissociates into

70
Q

van’t Hoff factor (i)

A

The number of ions that form when the compound dissociates

*Molecules that don’t ionize: i = 1

71
Q

Formula for dilution questions

A

Dilution factor = Final volume/Initial volume

20 = 50 mL/x
Solve for x –> x =2.5 mL

72
Q

When to use Avogadro’s #

A

6.22 x 10^-23

73
Q

Ideal Gas Law Example Problem

A
74
Q

Calculating osmotic pressure of a solution

A

Osmotic pressure = iMR*T

R is gas constant
T is temp in K
M is molarity of solution
i is van’t hoff factor; how many ions dissociate in the compound

In the question, multiply each compound’s M by how many ions the compound dissociates into. The largest number has the greatest osmo pressure.

Malika Gen Chem MCAT (3 decks)

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