Ch. 2: The Periodic Table (Complete)

Question 1

CH 2.3

Question 2

what are the three key rules that control how valence electrons work in an atom?

Answer 2

1. L to R across a period = electrons and protons are added one at a time
2. Down the elements of a group = principle quantum number increases by one each time
3. elements can gain or lose electrons in order to achieve a stable octet formation representative of noble (inert) gases

Question 3

explain why: electrons and protons are added one at a time L to R across a period and what the effect is

Answer 3

1. as the positivity of the nucleus increases, the electrons surrounding the nucleus experience a stronger electrostatic pull toward the atom’s center
2. this causes the electron cloud (the outer boundary defined by the valence shell electrons) moves closer and binds more tightly to the nucleus

Question 4

defn: effective nuclear charge (Zeff)

what is the effect of nonvalence electrons on Zeff

Answer 4

a measure of the net positive charge experienced by the outermost electrons in the electrostatic attraction between the valence shell electrons and the nucleus

the pull is somewhat mitigated by nonvalence electrons residing closer to the nucleus

Question 5

how does Zeff change from L to R for elements in the same period?

Answer 5

Zeff increases

Question 6

what are the implications of the principal quantum number increasing down the elements of a given group? (5)

Answer 6

1. the valence electrons are increasingly separated from the nucleus by a greater number of filled principal energy levels (inner shells)
2. this means a reduction in the electrostatic attraction between the valence electrons and the positively charged nucleus
3. the outermost electrons are held less tightly
4. increased shielding created by inner shell electrons cancels the increased positivity of the nucleus
5. thus Zeff is pretty constant among the elements within a group

Question 7

defn: octet rule

Answer 7

elements (esp. the ones with biological roles) tend to be most stable with 8 electrons in their valence shell, although there are many exceptions to this rule

Question 8

basis + defn: atomic radius

Answer 8

basis: imagine an atom as a cloud of electrons surrounding a dense core of protons and neutrons

atomic radius = one-half of the distance between the centers of two atoms of an element that are briefly in contact with each other

Question 9

what happens to atomic radius L to R across a period and why (4)?

Answer 9

atomic radius DECREASES L to R across a period

1. protons and electrons are added one at a time to the atoms
2. electrons are only added to the outermost shell, the number of inner-shell electrons remains constant
3. this increases the positive charge of the nucleus
4. which then pulls the outer electrons more closely inward and holds them more tightly

Question 10

what happens to atomic radius down a group and why (4)?

Answer 10

atomic radius INCREASES down a group

1. increasing principal quantum number down a group –> 2. valence electrons will be found further from nucleus 3. bc the # of inner shell electrons is increasing –> 4. separating the valence shell from nucleus

Question 11

what two generalizations must be made to understand ionic radii?

Answer 11

1. metals lose electrons and become positive; nonmetals gain electrons and become negative
2. metalloids can go in either direction, but tend to follow the trend based on which side of the metalloid line they fall on

Question 12

what is the trend (2) in ionic radii for nonmetals close to the metalloid line in comparison to nonmetals close to group VIIIA?

Answer 12

nonmetals close to the metalloid line:
1. require more electrons than other nonmetals to achieve octet (gain electrons)
2. have LARGER ionic radius than counterparts

Question 13

what is the trend in ionic radii for metals close to the metalloid line in comparison to metals close to group VIIIA?

Answer 13

metals close to the metalloid line:
1. have more electrons to lose to achieve octet (lose electrons)
2. have much SMALLER ionic radii than other metals

Question 14

defn and aka: ionization energy

Answer 14

aka: ionization potential

the energy required to remove an electron from a gaseous species

Question 15

explain the process associated with “ionization energy”. is this an endothermic or exothermic process? why?

Answer 15

ionization energy = energy required to remove an electron from a gaseous species (SO PROCESS = IONIZATION)

this is an ENDOTHERMIC process because removing an electron from an atom always requires heat input

Question 16

how does Zeff impact ionization energy?

what does this mean in terms of the trend of ionization energy on the periodic table?

Answer 16

1. the greater the atom’s Zeff (the closer the valence e’s are to the nucleus) –> 2. the more bound they are –> 3. the more difficult to remove one or more electrons –> 4. increases ionization energy

Ionization energy INCREASES from L to R across a period and from bottom to top in a group

Question 17

expln + defn: first ionization energy and second ionization energy

Answer 17

the subsequent removal of a second or third electron requires increasing amounts of energy because the removal of more than one electron means that electrons are being removed from an increasingly cationic (positive species)

first ionization energy = the energy necessary to remove the first electron

second ionization energy = the energy necessary to remove the second electron from the univalent cation (X+) to form the divalent cation (X2+)

Question 18

which is larger: the first or second ionization energy?

Answer 18

the second ionization energy

Question 19

groups + natural state + expln for name: active metals

Answer 19

elements in groups 1 and 2

called active metals bc they have such low ionization energies

do not naturally exist in neutral forms, but are always found in ionic compounds, minerals, or ores

Question 20

what changes must be made to group IA (alkali metals) or group IIA (alkaline earth metals) to form a stable, filled valence shell?

Answer 20

group IA –> the loss of one electron

group IIA –> the loss of 2 electrons

Question 21

why are the second ionization energies for group 1 monovalent cations disproportionally larger than the ones for group 2 or other monovalent cations?

Answer 21

removing one electron from a group 1 metal results in a noble gas-like electron configuration

Question 22

why are the ionization energies of nobel gases so high?

Answer 22

they are the least likely to give up electrons, have a stable electron configuration already

Question 23

what is the periodic trend of ionization energy?

Answer 23

L to R increases

bottom to top INCREASES

Question 24

defn: electron affinity

Answer 24

the energy dissipated by a gaseous species when it gains an electron

the opposite concept of ionization energy

Question 25

is the process associated with electron affinity endothermic or exothermic?

Answer 25

exothermic –> gaining an electron expels energy in the form of heat

deltaHrxn has a negative sign, but electron affinity is reported as positive bc it refers to the energy dissipated

Question 26

why are halogens the most greedy elemental group when it comes to electrons?

Answer 26

they only require one additional electron to complete their octet

Question 27

what is the periodic trend for electron affinity? why (@)?

Answer 27

L to R = increases –> why? 1. the stronger the electrostatic pull (the higher the Zeff) between the nucleus and the valence shell electrons –> 2. the greater the energy released will be when the atom gains an electron

top to bottom = decreases –> why? because the valence shell is farther away from the nucleus as principal quantum number increases

Question 28

describe the electron affinities of noble gases

Answer 28

EA on the order of 0 because they already possess a stable octet and cannot readily accept an electron

Question 29

defn: electronegativity

Answer 29

a measure of the attractive force that an atom will exert on an electron in a chemical bond

the greater the EN of an atom, the more it attracts electrons within a bond

Question 30

how is electronegativity related to ionization energy?

Answer 30

they adjust in parallel
- the lower the ionization energy, the lower the electronegativity

• the higher the ionization energy, the higher the electronegativity

Question 31

what three elements have negligible electronegativity? why?

Answer 31

the first three noble gases

because they do not form bonds often

Question 32

range: Pauling electronegativity scale

Answer 32

ranges from 0.7 to 4.0

Question 33

what is the trend of electronegativity across the periodic table?

Answer 33

L to R: increases

Top to bottom: decreases

Question 34

SUMMARIZE: periodic trends

Answer 34

Left to Right: atomic radius decreases // ionization energy, electron affinity, and electronegativity increases

Top to bottom: atomic radius increases // ionization energy, electron affinity, and electronegativity decreases

Question 35

CH 2.1

Question 36

defn: periodic law

Answer 36

the chemical and physical properties of the elements are dependent, in a periodic way, upon their atomic numbers (number of protons)

Question 37

defn: period vs. group/family

Answer 37

periods = rows
groups/families = columns

based on atomic number

Question 38

how do protons and electrons change across a period?

Answer 38

each element in a given period has one more proton and one more electron than the element to its left (in their neutral states)

Question 39

what is a commonality between elements within in a group (column) of the periodic table?

Answer 39

1. same electron configuration in their valence shell
2. share similar chemical properties

Question 40

what does the Roman numeral above each group represent?

Answer 40

the number of valence electrons elements in that group have in their neutral state

Question 41

defn: A elements vs. B elements

Answer 41

A elements = representative elements = groups IA through VIIIA
- have their valence electrons in the orbitals of either s or p subshells
–> electron configurations determined by the Roman numeral and the letter designation

B elements = nonrepresentative elements = includes transition elements (valence electrons in s and d subshells) and the lanthanide and actinide series (valence electrons in the s and f subshells)
–> may have unexpected electron configurations

Question 42

CH 2.2

Question 43

periodic place + what does this group include + char(6): metals

Answer 43

found on the left and middle of periodic table (includes active metals, transition metals, lanthanide and actinide series)
1. lustrous (shiny) (other than mercury)
2. high melting points and densities (other than lithium)
3. malleability
4. ductility (its ability to be pulled or drawn into wires)
5. easily gives up electrons
6. good conductors of heat and electricity (bc the valence electrons of all metals are only loosely held to their atoms, they are free to move)

Question 44

what are 2 unique characteristics of transition metals (many of the group B elements)?

Answer 44

have two or more oxidation states (charges when forming bonds with other atoms)

some transition metals are relatively nonreactive

Question 45

char at the atomic level (6): metals

Answer 45

1. low effective nuclear charge
2. low electronegativity/high electropositivity
3. large atomic radius
4. small ionic radius
5. low ionization energy
6. low electron affinity

Question 46

what subshells are the valence electrons of different metals in?

Answer 46

1. active metals = s subshell
2. transition metals = s and d subshells
3. lanthanide and actinide series = s and f subshells

Question 47

periodic place + char (5): nonmetals

Answer 47

predominantly on the upper right side of the periodic table
1. brittle when solid
2. show little to no metallic luster
3. poor conductors of heat and electricity
4. inability to easily give up electrons
5. less unified in their chemical and physical properties than the metals

Question 48

char at the atomic level (5): nonmetals

Answer 48

1. high ionization energies
2. high electron affinities
3. high electronegativities
4. small atomic radii
5. large ionic radii

Question 49

periodic place + aka + char (3): metalloids

Answer 49

separate the metals and nonmetals as a stair-step group of elements

aka = semimetals bc they share some characteristics with both metals and nonmetals

1. densities, melting points, and boiling points vary widely and can be combos of metallic and nonmetallic characteristics
2. reactivities depend on the elements with which they are reacting
3. good semiconductors bc of partial conductivity of electricity

Question 50

char at atomic level: metalloids

Answer 50

electronegativities and ionization energies lie between metals and nonmetals

Question 51

what are the 8 metalloids that form the staircase?

Answer 51

1. boron (B)
2. silicon (Si)
3. germanium (Ge)
4. arsenic (As)
5. antimony (Sb)
6. tellurium (Te)
7. polonium (Po)
8. astatine (At)

Question 52

diagram: periodic table, coded by element type

Answer 52

Question 53

CH 2.4

Question 54

group + char (6): alkali metals

Answer 54

Group IA/Group 1
1. have most of the classic physical properties of metals except their densities are lower than those of other metals
2. have only one loosely bound electron in their outermost shells
3. have low Zeff values
4. have the largest atomic radii of all elements in their rows
5. low ionization energy, low electron affinity, low electronegativity
6. easily lose one electron to form univalent cations and react readily with nonmetals (especially the halogens)

Question 55

group + char (4): alkaline earth metals

Answer 55

group IIA/group 2
1. possess many properties characteristic of metals
2. slightly higher Zeff than alkali metals
3. slightly smaller atomic radii than alkali metals
4. have 2 electrons in their valence shell, which are both easily removed to form divalent cations

Question 56

what are alkali and alkaline earth metals called together? why?

Answer 56

active metals = they are so reactive that they are not naturally found in their elemental (neutral) state

Question 57

group + char (6): chalcogens

Answer 57

group VIA/group 16
1. eclectic group of nonmetals and metalloids
2. crucial for normal biological functions
3. each have 6 electrons in their valence shell
4. have small atomic radii
5. large ionic radii
6. at high concentrations, many of these elements can be toxic or damaging, no matter how biologically useful

Question 58

what are the 3 biologically important elements in the chalcogen group? what are 2 characteristics of the rest of the elements in the group?

Answer 58

1. oxygen
2. sulfur
3. selenium

rest of the group: primarily metallic and generally toxic to living organisms

Question 59

group + char (8): halogens

what element in this group has the highest electronegativity of all elements?

Answer 59

group VIIA/group 17
1. highly reactive nonmetals with 7 valence electrons
2. desperate to complete their octets by gaining one additional electrons
3. variable physical properties
4. more uniform chemical reactivity
5. high electronegativity
6. high electron affinity
7. very reactive toward alkali and alkaline earth metals
8. so reactive that they are not found in their elemental state, but rather as ions (called halides)

Fluorine has the highest EN of all elements

Question 60

group + char (5) + aka: noble gases

Answer 60

group VIIIA/group 8
aka inert gases

1. minimal chemical reactivity due to their filled valence shells
2. high ionization energies
3. little or no tendency to gain or lose electrons
4. no measurable EN (for at least He, Ne, and Ar)
5. extremely low boiling points
6. exist as gases at room temperature

Question 61

group + char (8): transition metals

Answer 61

groups IB to VIIIB/groups 3 - 12
1. considered to be metals
2. low electron affinities, low ionization energies, low electronegativities
3. very hard
4. high melting and boiling points
5. malleable
6. good conductors due to loosely held electrons that progressively fill the d-orbitals in their valence shells
7. many of them can have different possible charged forms/oxidation states bc they are capable of losing different numbers of electrons from the s and d orbitals in their valence shells –> form many different ionic compounds —> different oxidation states often correspond to different colors (solutions with transition metal-containing complexes are often vibrant)
8. tend to associate in solution with molecules of water (hydration complexes) or with nonmetals –> this ability to form complexes adds to the variable solubility of certain transition metal-containing compounds

Question 62

explain why the formation of complexes allows for the vibrant colors characteristic of transition metals (3)

Answer 62

1. the formation of complexes causes the d-orbitals to split into 2 energy sublevels
2. this enables many of the complexes to absorb certain frequencies of light (those containing the precise amount of energy required to raise electrons from the lower to the higher energy d orbitals)
3. the frequencies not absorbed (subtraction frequencies) give the complexes their characteristic colors

Question 63

when we perceive an object as a particular color, it is because that color is not absorbed, but rather reflected, by the object. explain.

Answer 63

if an object absorbs a given color of light and reflects all others, our brain mixes these subtraction frequencies and we perceive the complementary color of the frequency that was absorbed