Unit 3.1 The Periodic Table & Periodic Trends

3.1.1 The Periodic Table

Structure of the Periodic Table

• The Periodic Table = A list of all known elements arranged in order of increasing atomic number (from 1 to 118)

• Elements are in rows and columns

  • Atoms with the same # of shells are placed in the same row

• Period = A row of elements

  • n = Period # = The outer energy level that is occupied by electrons

• Group = Elements aligned vertically in columns

  • These elements share a similar outer-shell electron configuration

  • Valence Electrons = electrons in the outer shell

The Blocks of the Periodic Table

All elements are categorized into 4 main blocks

• s-block = elements with only s electrons in the outer shell

• p-block = elements with at least 1 p-electron in the outer shell

• d-block = elements with at least 1 d-electron and at least 1 s-electron but NO f or p electrons in the outer shell

• f-block = elements with at least 1 f-electron and at least 1 s-electron but NO d or p electrons in the outer shell

Deducing Electron Configurations

• You can use the position of an element in the periodic table to help you figure out its electron configuration

• ex. Germanium = [Ar] 3d¹⁰4s²4p²

  • Ge is in Group 4 → Tells you that there are 4 Valence Electrons

  • Ge is in Period 4 → Tells you that the 4 VE are in the 4th shell

  • The 2nd position in the p-block → Tells you that 2 electrons are in the p-subshell

3.1.2 Periodic Trends: Physical - Atomic & Ionic Radius

Atomic & Ionic Radius

Atomic Radius

• Atomic Radius= the size of an atom = the distance between the nucleus of an atom and the outermost electron shell

  • Decreases across each period

  • Increases down each group

Electron Shell Theory

• As you move across a period, the amount of protons increases, thus increasing the pull of the nuclei on the electrons, so the atoms become smaller

• As you move down a group, more electrons are being added, so the radii increase

  • However, the outer electrons don’t feel the pull of the positive nuclei charge as much because the inner electrons repel them

Ionic Radius

• Ionic Radius = the size of an ion

  • Trend down a group is the same as atomic radius: as the number of shells increases, the radius increases

  • However, the trend across a period depends on whether an ion is positive or negative

• In general:

  • As the negative charge increases, the ionic radii increase

  • As the positive charge increases, the ionic radii decrease

Electron Shell Theory

• Negative ions are formed when an atom accepts extra electrons which experience repulsion between the other valence electrons

  • So that causes a larger ionic radius

• Positive ions are formed when an atom loses electrons

  • Since there are fewer electrons while the nuclear charge remains the same, each remaining electron feels the pull of the nucleus more strongly than before

  • Thus, the ionic radius decreases

3.1.3 Periodic Trends: Physical - Ionization Energy

First Ionization Energy

• Ionization Energy (IE) = The amount of energy required to remove 1 mole of electrons from 1 mole of atoms (in the gaseous state) to form 1 mole of gaseous ions

  • Measured in kilojoules per mole (kJ mol^-1)

• First Ionization Energy = the amount of energy to remove the first mole of electrons

Ionization Energies: Trends

• Increases across a period

• Decreases down a group

Trends influenced by 4 factors:

1. Size of the Nuclear Charge

   • As the atomic number (# of protons) increases

   • → the nuclear charge increases

   • → greater attractive forces between positive nucleus and outer electrons

   • → more energy is required to overcome these strong attractive forces when trying to remove an outer electron

2. Distance of Outer Electrons from the Nucleus

   • Electrons in shells further from the nucleus feel its pull less

   • → So the further out electrons are, the lower ionization energy they have

3. Shielding Effect of Inner Electrons

   • = When the electrons in the inner shells repel those in the outer shells, so they don’t feel the attractive forces of the nucleus as much

   • The greater the shielding, the lower the ionization energy of the outer electrons

4. Spin-Pair Repulsion

   • Paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals, so it’s easier to remove an electron

   • Tip: This is the reason why the first ionization energy is always the lowest

Ionization Energy (IE) Across a Period

• IE increases across a period because:

  • The nuclear charge across a period increases

  • The distance between the nucleus and outer electrons is the same

  • The shielding of the outer electrons by the inner is the same

Exceptions

• IE decreases suddenly between the last element in one period and the first element in the next period because:

  • The distance between the nucleus and the outer electrons increases

  • The inner electrons’ shielding increases

  • These 2 factors outweigh the increased nuclear charge

• IE slightly decreases between beryllium and boron since the 5th electron in boron is in the 2p subshell (which is further away from the nucleus than the 2s subshell of beryllium)

  • beryllium = 900 kJ mol^-1

  • boron = 801 kJ mol^-1

• IE slightly decreases between nitrogen and oxygen because there is spin-pair repulsion in the 2p subshell of oxygen

  • nitrogen = 1402 kJ mol^-1

  • oxygen = 1314 kJ mol^-1

Ionization Energy (IE) Down a Group

• As you go down a group, IE decreases because

  • The distance between the nucleus and the outer electrons increases

  • Electron shielding increases

  • The effective nuclear charge decreases as shielding increases

Successive Ionization Energies of an Element

• Successive ionization energies (the amount of energy when you keep taking off electrons) increases because removing the next electron always takes more energy than the previous

  • In other words, it is harder to remove an electron from an already positive ion than a neutral atom

• For several reasons:

  • As you remove more electrons, the shielding decreases, so the attractive forces increase along with an increased proton to electron ratio (so each electron feels the pull of the nucleus more and more)

• Note: increases in successive ionization energy is NOT constant, but dependent on each atom’s electron configuration

3.1.4 Periodic Trends: Physical - Electron Affinity

• Anions = negative ions formed when electrons gain electrons

• Electron Affinity (EA) = The amount of energy released when 1 mole of electrons is gained by 1 mole of atoms of an element in the gaseous state to form 1 mole of gaseous ions

  • Think of it as the opposite of Ionization Energy

  • Measured in kilojoules per mole (kJ mol^-1)

• The periodicity of Electron Affinity has a similar pattern to Ionization Energies (but inverted)

• The strongest “pull” on electrons → The greater amount of energy released when negative ions are formed

• EA decreases down a group because:

  • As the atoms take on more electrons, the attraction for more decreases because the shielding increases and the effective nuclear charge decreases

3.1.5 Periodic Trends: Physical - Electronegativity

Electronegativity: Definition

• Electronegativity = The ability of an atom to attract an electron pair in a covalent bond

  • Caused by the positive nucleus being able to pull the negative electrons closer

  • The Pauling Scale = A way to measure the value of electronegativity for each atom (on an arbitrary scale from 0.0 to 4.0)

Electronegativity: Factors

Nuclear Charge

• The more positively charged protons in the nucleus, the more nuclear attraction between them and the negatively charged electrons

  • Thus, an increase in nuclear charge means an increase in electronegativity

Atomic Radius

• Atomic Radius = The distance between the nucleus and valence electrons

• As the atomic radius increases, the valence electrons are farther away, so they have a decreased attraction to the nucleus and less electronegativity

Electronegativity: Trends

Down a Group

• Electronegativity decreases down a group

• Although more protons are added and you would expect the electronegativity to increase with the increased nuclear attraction, this isn’t the case due to electron shielding:

  • As more electrons are being added, more shells are being filled, so more shells means more distance between the nucleus and valence electrons

  • In this case, the increased distance (that causes a decreased electronegativity) matters more than the increased nuclear charge (that would have caused an increase in electronegativity)

Across a Period

• Electronegativity increases across a period

• As you go across a period, the number of protons increases, while the shielding remains the same across a period, since no new shells are being added

  • Therefore, the increased nuclear charge matters more in this case

• So the increased nuclear charge causes

  • Decrease in atomic radius (pulls electrons closer to nucleus)

  • Increase in attraction between nucleus and valence electrons

  • Increase in Electronegativity

3.1.6 Periodic Trends: Chemical

Metallic & Non-metallic Properties

• What causes these different properties? Trends in:

  • Atomic Radius

  • Ionic Radius

  • Electron Affinity (EA)

  • Ionization Energy (IE)

  • Electronegativity (EN)

• Low EN and low IE of metals allow their valence electrons to be able to move away from the nucleus or “delocalize”

• High EN and high EA of non-metals allows for their electrons to be shared and form covalent bonds

• Similarities in EN for the elements in the diagonal band separating metals from non-metals explains their behavior as" “metalloids”

Unit 3.2 Oxides, Group 1 & Group 17

3.2.1 Periodic Trends: Oxides Across a Period

Oxides Across a Period

• Oxide = a binary compound that contains oxygen and another element

  • ex. Carbon dioxide CO₂

• Since oxides are acid-base they show their chemical trend: they change from basic to amphoteric to acidic as you move across a period

• Amphoteric = having the ability to react chemically as either an acid or a base

• ex. Aluminum oxide (so it can react with an acid like HCl or with a base like NaOH)

Period 3 Oxides

• The reason for the different acidic or basic natures is because of their structure, bonding, and electronegativity

• Electrons will be transferred to oxygen when forming oxides and providing an ionic bond because of the highest difference in electronegativity between oxygen and Na/Mg/Al

• On the other hand, Si, P, and S (elements with the least difference in EN) will share the electrons with oxygen to form covalently bonded oxides

3.2.2 Periodic Trends: Group 1 - The Alkali Metals

• Group 1 Metals = “Alkali Metals” = metals that form alkaline solutions with high pH when they react with water

• All elements in group 1 end with the electron configuration: ns¹

Physical Properties

• Soft and easy to cut (more so as you move down the group)

• Shiny

• Conduct heat/electricity

• Low melting points (decreases more going down the group since atomic radius increases so metallic bonding gets weaker)

• Low Density

Chemical Properties

• Reactive with oxygen and water in the air

  • Often kept under oil to prevent reactions

• Highly reactive with water

  • And produces an alkaline metal hydroxide solution and hydrogen gas

3.2.3 Periodic Trends: Group 17 - The Halogens

The Halogens

• Group 17 non-metals

• Poisonous

• Diatomic = forming molecules of 2 atoms

• Have 7 valence electrons

• Form halide ions by gaining one more electron to complete the octet rule

Physical Properties Trends (Down the Group)

• Density, Melting/Boiling Points of the halogens increase going down the group

• Reactivity decreases down the group

  • Why?

    • Halogens electron configurations all end in ns²np⁵

    • When the 7 valence electrons react, they need one more to have a full shell

    • Electron Affinity decreases

    • Atomic Radius increases

      • → # of Shells increases

      • → Shielding increases

      • → Distance to nucleus increases

      • → Weaker electrostatic forces to attract the 8th electron

      • Therefore, as you go down the group, the reactivity decreases (because it is harder to attract the 8th electron)

Reaction of Halogens with Halid Ions in Displacement Reactions

• Halogen Displacement = When a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide

Halogen Displacement Reactions

Chlorine and Bromine

• Chlorine + colorless potassium bromide = bromine (Orange Solution)

• Since chlorine is above bromine in group 17, it is more reactive

  • So chlorine will displace bromine

2KBr (aq)       +    Cl₂ (aq)  →     2KCl (aq)     +       Br₂(aq)

potassium bromide + chlorine → potassium chloride + bromine

Bromine and Iodine

• Since bromine is higher than iodine in group 17, it is more reactive

• Thus, bromine will displace iodine

  Br₂ (l) +      2NaI (aq)    →    2NaBr (aq)     +    I₂ (aq)

bromine +  sodium Iodide → sodium bromide + iodine