Periodic Trends

Question 1

First ionization energy

Answer 1

• The amount of energy needed to remove the first electron from a neutral atom (the outermost electron)
• In general, first ionization energy increases going up and right in the periodic table

Question 2

Effective nuclear charge and ionization energy

Answer 2

Elements with a higher effective nuclear charge have a higher first ionization energy because of the increased attraction from the protons making it more difficult to remove electrons from the atom (“ionize them”)

Question 3

Ionization energy trend:

Answer 3

Increases going to the right of the periodic table and going up the periodic table (noble gases are included)

Question 4

Elements with the _____ first ionization energy are the ones that will most readily lose an electron and become ionized

Answer 4

LOWEST
Ex: potassium has a lower ionization energy than lithium and therefore will most readily lose an electron

Question 5

Ionization energy

Answer 5

The amount of energy required to remove an electron from an atom. The easier it is to remove an electron, the lower the ionization energy

Question 6

General trend form atomic radius for atoms with a neutral charge

Answer 6

Increases going down and to the left of the periodic table

Atoms get larger going down: due to higher energy “n” shell = larger size
Atoms get smaller going across: due to higher effective nuclear charge = smaller size

Question 7

General trend for atomic radius for isoelectronic ions

Answer 7

More positive charge = smaller radius
More protons = more attraction to electrons = smaller size atom

More negative charge = larger radius
Less protons = less attraction to electrons = larger size atom

Question 8

What are the two factors that affect the size of an ion?

Answer 8

The number of atomic orbitals: an ion with more atomic orbitals will have a larger atomic radius because these atomic orbitals occupy more space and expand the electron cloud

The effective nuclear charge: an ion with fewer protons will have a larger atomic radius because protons pull in the electron cloud, reducing its range. The fewer protons in the nucleus, the weaker the pull on electrons

Question 9

What is the reason for atoms size increasing going down a column on the periodic table?

Answer 9

This is due to electron shielding

The electron shielding effect reduces the effect of the nuclear charge. As more core electrons are added, they begin to shield the valence electrons from the increasingly positive nucleus which allows the outer electrons to move further away from the nucleus and therefore increasing the atoms size

Question 10

Electron affinity

Answer 10

The amount of energy released when an atom gains an electron

Question 11

Electronegativity

Answer 11

• The ability of an atom to attract electrons in a bond
• The higher the electronegativity of an atom, the greater ability it has to attract an electron pair

Question 12

Which atoms would most readily form ionic compounds?

Answer 12

The ones with the lowest ionization energy (furthest to the left and down a column) because they will most readily lose electrons to become ionized

Question 13

Noble gases exert what type of force?

Answer 13

Very weak london dispersion forces (and therefore have very low melting and boiling points)

Question 14

Characteristics of halogens:

Answer 14

7 valence electrons
High electronegativity
High electron affinity
Easily reduced (gain one electron to have a full octet)
Strong oxidizing agents
Highly reactive with metals

Question 15

properties of metals

Answer 15

• malleable, lustrous
• good conductors of heat/electricity
• form basic oxides
• lose electrons to form cations
• usually solid at room temperature (except Hg-liquid)
• high melting and boiling points

Question 16

properties of non-metals

Answer 16

• brittle, dull
• poor conductors of electricity/heat
• form acidic oxides
• gain electrons to form anions
• gas or solid at room temperature (except Br-liquid)
• low melting and boiling points

Question 17

trend when comparing melting points of compounds

Answer 17

comparing their ionic character: difference in their electronegative –> the larger the difference in electronegativity, the greater their ionic character and the higher the melting point

Question 18

trend for metallic character:

Answer 18

increases going from right to left and going down a group

Question 19

trend for electronegavity:

Answer 19

• increases going to the right and up the periodic table (excluding the noble gases
• F is the most EN

noble gases have full valence shells –> no EN value

Question 20

what is the reasoning for the EN trend?

Answer 20

• increasing EN from left to right: increasing number of protons increases ability of an atom to attract an electron pair
• decreasing EN going down a group: as atomic radius increase, the valence electrons experience greater electron sheilding and the ability of an atom to attract an electron pair is decreased

Question 21

electron affinity trend:

Answer 21

• increases going to the right and up on the periodic table (excludes noble gases)

Question 22

reasoning for electron affinity trend:

Answer 22

• increas in electron affinity going to the right: greater nuclear attraction between protons and electrons causes a stronger affinity for electrons
• decrease in electron affinity going down a group: attraction of electrons to the nucleus decreases due to electron sheilding causing a decrease in electron affinity

Question 23

what atoms are exceptions in the electron affinity trend?

Answer 23

• Phosphorus and silicon: even though P is to the right of S, P has a half-filled subshell electron configuration –> S has a higher electron affinity because gaining an electron will give it the half filled subshell configuration (more stable)

Question 24

difference between electronegativity and electron affinity?

Answer 24

electronegativity is specific to electron pairs in a bond whereas electron affinity is the addition os a single electron

Question 25

ionization energy of alkaline earth metals and group 13 exception:

Answer 25

• alkaline earth metals (group 2) have a greater ionization energy than group 13 elements
• alkaline earth metals have full filled orbitals and require more energy to remove an electron than group 13 metals

Question 26

ionization energy of group 15 and group 16 exception:

Answer 26

• group 15 elements have higher ionization energy than group 16 elements
• group 15 have half-filled orbitals, which is a more stable configuration and requires more energy to remove an electron

Question 27

ionic bonds

Answer 27

• complete transfer of electrons between 2 atoms
• metal loses electrons to become a positively charged cation
• non-metal accepts the electron and becomes a negatively charged anion

Question 28

sigma bonds

Answer 28

• covalent bonds made by end-to-end overlap of atomic orbitals
• stronger than pi bonds

Question 29

pi bonds

Answer 29

• covalent bonds made by side-to-side (lateral) overlap of p-orbitals