Gen Chem Terms

Question 1

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

Answer 1

atoms, charge

Question 2

Is there an error in the balancing of this reaction?

Answer 2

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

Question 3

Calculating theoretical yield question

Answer 3

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.

Question 4

Calculating percent yield

Answer 4

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

Question 5

Bronsted Lowry acid vs Lewis Acid definitions

Answer 5

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

Question 6

Question about BL acid and Lewis acid

Answer 6

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.

Question 7

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

Answer 7

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

Question 8

Difference between noncovalent IMFs

Answer 8

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

Question 9

Dipole-induced dipole IMF

Answer 9

Question 10

Properties of metals

Answer 10

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

Question 11

Examples of coordinate covalent bonds

Answer 11

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

Question 12

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

Answer 12

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

Question 13

Metallic character increases as you go ______ the periodic table.

Answer 13

down

Question 14

Condosity

Answer 14

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

Question 15

Ionic compound exceptions

Answer 15

NH4Cl, (NH4)2SO4, NH4BF4

Question 16

Ranking compounds based on ionic character

Answer 16

-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

Question 17

Characteristics of covalent bonds

Answer 17

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

Question 18

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

Answer 18

greater than 1.7

Question 19

What is meant by the most stable electron configuration?

Answer 19

The electron configuration of the atom in its ground state.

Question 20

gamma radioactive decay

Answer 20

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

Question 21

alpha radioactive decay

Answer 21

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

Question 22

B+ decay (positron emission)

Answer 22

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

Question 23

Beta decay: electron capture

Answer 23

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

Question 24

B- decay (electron emission)

Answer 24

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

Question 25

what happens molecularly during fluorescence

Answer 25

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).

Question 26

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

Answer 26

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.

Question 27

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

Answer 27

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

Question 28

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

Answer 28

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

Question 29

Compare insulators & conductors

Answer 29

Question 30

Flashpoint

Answer 30

The temp at which a metal decomposes or explodes

Question 31

What are the subshells?

Answer 31

s, p, d, f

Question 32

Pauli exclusion principle

Answer 32

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

Question 33

Aufbau principle

Answer 33

Electrons must be filled from lower energy to higher energy subshells

Question 34

What are orbitals?

Answer 34

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

Question 35

principle (n) quantum #

Answer 35

principle quantum number, describes the number of shells

Question 36

azimuthal (l) quantum #

Answer 36

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

Question 37

How to calculate l

Answer 37

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

Question 38

magnetic (ml) quantum #

Answer 38

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

Question 39

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

Answer 39

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

Question 40

Hund's rule

Answer 40

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

Question 41

Notation for writing an element from the periodic table

Answer 41

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

Question 42

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

Answer 42

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

Question 43

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

Answer 43

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

Question 44

Electronegativity definition & trend

Answer 44

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.

Question 45

Ionization energy definition & trend

Answer 45

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.

Question 46

Answer 46

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)

Question 47

Electron affinity definition & trend

Answer 47

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.

Question 48

Exceptions to electron affinity

Answer 48

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.

Question 49

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

Answer 49

True

Question 50

Ionic radius definition & trend

Answer 50

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.

Question 51

Effective nuclear charge definition and trend

Answer 51

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.

Question 52

Calculating effective nuclear charge

Answer 52

protons - # of innershell core electrons (nonvalence)

Question 53

Atomic radius definition & trend

Answer 53

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.

Question 54

How does effective nuclear charge affect atomic radius?

Answer 54

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.

Question 55

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

Answer 55

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

Question 56

What is the oxidizing agent in this reaction?

Answer 56

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.

Question 57

Hard question

Answer 57

Question 58

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

Answer 58

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

Question 59

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

Answer 59

True, ionization energy and reactivity are inversely related.

Question 60

Properties of alkaline earth metals

Answer 60

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

Question 61

Properties of alkali metals

Answer 61

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

Question 62

Relationship between atomic radius & strength of bond

Answer 62

Smaller atomic radius = stronger bond = less acidic

Question 63

Calculating mass percent of a component in a mixture/substance

Answer 63

mass percent = mass of component/total mass of mixture

Question 64

hard question

Answer 64

Question 65

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

Answer 65

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.

Question 66

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

Answer 66

180˚, 120˚, and 109.5˚

Question 67

Ideal gas assumptions:

Answer 67

-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

Question 68

Deviations from ideal gas

Answer 68

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

Question 69

Calculating osmolarity of a solution

Answer 69

molarity of solution x n n is the # of ions that the solution dissociates into

Question 70

van't Hoff factor (i)

Answer 70

The number of ions that form when the compound dissociates *Molecules that don't ionize: i = 1

Question 71

Formula for dilution questions

Answer 71

Dilution factor = Final volume/Initial volume 20 = 50 mL/x Solve for x --> x =2.5 mL

Question 72

When to use Avogadro's #

Answer 72

6.22 x 10^-23

Question 73

Ideal Gas Law Example Problem

Question 74

Calculating osmotic pressure of a solution

Answer 74

Osmotic pressure = i*M*R*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.