ch 2 - periodic properties

Question 1

group 1 elements

Answer 1

alkali metals
- excluding hydrogen

Question 2

group 2 elements

Answer 2

alkaline earth metals

Question 3

group 3-12 elements

Answer 3

transition metals

Question 4

groups 13-17 elements

Answer 4

metalloids
- has a combination of both metaallic nd non-metallic characteristics

Question 5

group 17 elements

Answer 5

halogens

Question 6

group 18 elements

Answer 6

noble gases

Question 7

period 6 and 7 elements

Answer 7

inner transition metals
- separated from the transition metals because their properties differ

Question 8

period 6 elements

Answer 8

lanthanides

Question 9

period 7 elements

Answer 9

actinides

Question 10

what are the 7 diatomic atoms

Answer 10

hydrogen, nitrogen, fluorine, oxygen, iodine, chlorine, bromine
- Have No Fear Of Ice Cold Beer

Question 11

how does metallic character increase on the periodic table

Answer 11

going right to left across a period
- and going down a group

Question 12

characteristics of metals

Answer 12

malleable, lustrous
- good conductors of electricity, heat
- form basic oxides
- lose electrons to form cations
- usually solid at liquid temperature, with the exception of Hg (liquid)
- generally, high melting and boiling points

Question 13

characteristics of non-metals

Answer 13

brittle, dull
- poor conductors of electricity/heat
- form acid oxides
- gain electrons to form anions
- gas or solid at room temperature, with the exception of Br (liquid)
- generally, low melting and boiling points

Question 14

atomic radius: definition, trend, and reasoning

Answer 14

definition
- half the distance between the nuclei of two identical atoms bonded together

trend
- the atomic radius decreases from left to right across a period and increases going down a group

reasoning
- across a period, the number of protons in an atom increases
- increasing protons results in greater nuclear attraction between the protons and electrons, which results in shells being pulled closer to the nucleus (which equals a smaller radius)
- going down a group, the number of electrons increases
- each additional energy level gets farther and farther away from the nucleus, which causes the atomic radius to increase

Question 15

effective nuclear charge: definition, trend, reasoning

Answer 15

definition
- the effective nuclear charge (Zeff) is the nuclear charge experienced by an electron in an atom with multiple electrons
- this charge is assigned due to a shielding effect of electrons preventing other electrons in higher orbitals from experiencing a strong attraction to the nucleus
- this effect explains why valence electrons are more easily removed
- effective nuclear charge equation: Zeff = Z-S, where Z = # of protons, S = # of shielding electrons

trend
- effective nuclear charges increases across a period from left to right and decreases going down a group

reasoning
- across a period, the number of protons are increasing with no increase in a shielding effect, which results in electrons being pulled closer to the nucleus due to a stronger attraction
- going down a group, more shielding causes the effective nuclear charge to decrease
- as electrons get further away from the nucleus, the attractive force between protons and electrons naturally lessens

Question 16

what is the equation for effective nuclear charge?

Answer 16

Zeff = Z - S
- Z = number of protons
- S = number of shielding electrons

Question 17

shielding effect

Answer 17

electrons in higher orbitals do not have strong attraction to nucleus due to shielding effect of closer electrons
- explains why valence electrons are more easily removed
- the electrons between the electron of interest and nucleus cancels some of the positive nuclear charge → resulting in a weaker attraction

part of effective nuclear charge

Question 18

when a neutral atom loses an electron and becomes a cation, what happens to effective nuclear charge & radius?

Answer 18

the effective nuclear charge increases, which causes the radius to decrease
- this is because protons will pull in electrons closer to the nucleus

Question 19

when a neutral atom gains an electron and becomes an anion, what happens to effective nuclear charge & radius

Answer 19

atomic radius will increase, due to a decrease in effective nuclear charge
- causes a decrease in the pull from the protons in the nucleus

Question 20

do metals usually form anions or cations?

Answer 20

they typically form cations
- resulting in their ionic radius to be less than their atomic radius

Question 21

do non-metals usually form anions or cations?

Answer 21

they typically form anions
- resulting in their ionic radius to be greater than their atomic radius

Question 22

isoelectronic series: definition only

Answer 22

these are atoms that have the same electron configuration, but differing number of protons
- in this case, the most positively charged atom will have the smallest radius because it has the most protons with the same number of electrons
- therefore will have the greatest attraction to the nucleus

Question 23

do anions or cations have a larger radius

Answer 23

anions have a larger radius than cations

Question 24

Q: which of the following ions has the smallest radius?
a) Na+
b) Br-
c) S2-
d) P3-
e) Al3+

Answer 24

Al3+
- can eliminate anions as a potential answer because anions have a larger radius than cations
- when comparing Na+ and Al3+, you should be able to determine that these cations are isoelectronic
- however, since Al3+ has more protons, the electrons will experience a greater pull resulting in a smaller radius

Question 25

Q: what is the effective nuclear charge for a valence electron of Selenium?

Answer 25

Zeff of Se = +6 - the first step in determining effective nuclear charge is to find the number of shielding electrons (non-valence electrons) - looking at the periodic table, we can determine that Selenium has 6 valence electrons - it's non-valence electrons would be 34-6 = 28 - Se has 34 protons Zeff = Z - S Zeff = 34 - 28 = +6 note: the effective nuclear is equal to the number of valence electrons - this is often the case, but not always

Question 26

ionization energy: definition, trend, reasoning

Answer 26

definition - the energy needed to remove na electron from an atom trend - ionization energy increases going from left to right across a period and decreases down a group reasoning - as you go from left to right across the periodic table., the number of electrons and protons increases - as the valence shell continues to fill, the electrons become harder to remove (require more energy) due to an increase in effective nuclear charge e.g. fluorine has a high ionization energy because its electrons are strongly attracted to the more positively charged atomic nucleus

Question 27

as the valence shell continues to fill, do the electrons become easier or harder to remove?

Answer 27

electrons become harder to remove and require more energy - this is due to an increase in effective nuclear charge

Question 28

going down a group, does the shielding effect increase or decrease?

Answer 28

shielding effect increases - outer electrons are not experiencing the full pull of the positive charge of protons due to increasing inner electrons shielding them - this results in a lower ionization energy as you go down a group because electrons become easier to expel with increasing shells e.g. iodine - has a lower ionization energy than fluorine because it contains more electron shells, which makes it easier to extract an electron from the atom's valence shell

Question 29

multiple ionization energies

Answer 29

the first ionization energy is the energy required to remove the outermost electron - following removal of the first electron, elements can ave second, third, fourth, etc. ionization energies - these values are always larger than the first ionization energy because subsequent electrons are more difficult to remove e.g. sodium, the first ionization energy is 496 kJ/mol - however, following that, there is a huge jump to 4562 kJ/mol for the second ionization energy because we would be removing an electron from a stable configuration (a full outer shell)

Question 30

2 exceptions to the rules to ionization energy

Answer 30

1. alkaline earth metals (group 2) have greater ionization energy than group 13 elements - this is because alkaline earth metals have completely filled orbitals and require more energy to remove an electron than group 13 elements 2. group 15 elements have greater ionization energies than group 16 elements - group 15 elements have half-filled orbitals, which is a more stable configuration than that of group 16 elements and thus require more energy to remove an electron

Question 31

Q: which of the following elements has the highest ionization energy? a) sodium b) fluorine c) chorine d) carbon e) neon

Answer 31

e) neon - this is because it is a noble gas - it is essentially impossible to extract an electron from neon's valence hell because it is completely filled and therefore VERY stable - noble gasses are chemically inert because their high degree of stability makes it unlikely for them to lose or gain electrons

Question 32

electron affinity: definition, trend, reasoning

Answer 32

definition - the amount of energy released when an electron is added to an atom trend - electron affinity increases going from left to right across a period and decreases going down a group reasoning - going across a period, greater nuclear attraction between protons and electrons creates a stronger affinity for electrons - going down a group, the attraction of an electron to the nucleus decreases due to shielding - electron affinity, therefore decreases note: similar to ionization energy, the trend for electron affinity is a general trend and may not apply to all elements - e.g. although phosphorous is to the right of silicon, phosphorous has a half-filled subshell electron configuration - silicon has a greater affinity for an electron than phosphorous because accepting an electron will give it the half-filled subshell configuration

Question 33

electronegativity: definition, trend, reasoning

Answer 33

definition - the ability of an atom to attract electrons in a bond - the higher the electronegativity of an atom, the greater ability to attract an electron pair trend - electronegativity increases going from left to right across a period and decreases down a group - the most electronegative element is fluorine reasoning - with increasing protons as you go from left to right across. period, the ability of an atom to attract an electron pair is increased - this is similar to the electron affinity trend, however this trend is specific to electron pairs in a bond, not addition of a single electron - going down a group, as the atomic radius increases, and the valence electrons experience greater shielding, the ability to attract an electron pair is decreased - since fluorine is the most electronegative atom, it is clear to see that moving towards fluorine electronegativity must increase (left to right) and moving away from fluorine (top to bottom) electronegativity must decrease note: the noble gases are an exception to this trend as they have full valence shells and therefore NO electronegativity value