5. Periodicity // R3.2, S2.2-3

Alkali metals

1. Group 1

2. Metals

3. Very reactive — oxidation — Losing their outer electron to form the metal ion.

4. Good reducing agents

5. Reactivity increases down the group — outer electron is in higher energy levels and requires less energy to remove it

6. Stored under liquid paraffin to prevent them reacting with air

What do Alkali Metals react with

1. All react with H₂O → Metal hydroxide alkaline solution + H₂ (g)

2. React readily with Cl, Br, and I to form ionic salts

3. Li — floats and reacts quietly

4. Na — melts into a ball which darts around on the surface

5. K — heat generated from the reaction ignites hydrogen

Halogens

1. Group 17

2. Non metals

3. Very reactive (reduction) and replacement reactions

4. Good oxidizing agents

5. Create diatomic molecules

6. Reactivity decreases down the group (outer shell is at increasingly higher energy levels and further from the nucleus + more electrons between the nucleus and the outer shell, decreases the attraction for an extra electron)

How do Halogens react

By gaining one more electron to form halide ions

Chlorine is a stronger oxidizing agent than bromine, but both are stronger than Iodine:

How to test for Halide Ions

Add silver nitrate solution

Ag⁺ (aq) + X^- (aq) → AgX (s) —(Light)→ Ag (s) + (1/2)X₂

X = Cl, Br, I

Silver ions react with the halide ions to form a precipitate of the silver halide.

The silver halides can be distinguished by their colour.

Basis of old-fashioned film photography.

1. AgCl = White

2. AgBr = Cream

3. AgI = Yellow

Nobel Gasses

Group 18; non metals, not acids nor bases, low reactivity, monoatomic, weak intermolecular forces (LDP)

Diatomic

Molecules composed of only two atoms, of the same or different chemical elements. E.g. H₂ or O₂ or CO

Group

Column in the PT. Elements in the same group have same number of valence electrons

Period

Row in the PT. Periodic number explains the number of energy levels

Block

s,p,d,f. Indicates the last occupied electron orbital

Atomic Radius

Half the distance between the nuclei of two bonded atoms or the distance from the nucleus to the outermost electron.

Atomic Radius trend

Becomes bigge

r down a group and from right to left on a period

Factors affecting Atomic Radius

• Nuclear charge
  – Attractive force that pulls all the electrons closer to the nucleus. With an increase in nuclear charge, the atomic radius becomes smaller.

• Shielding effect by the electrons of the inner shell(s)
  – this makes the atomic radius larger. The shielding effect is the result of repulsion between the electrons in the inner shell and those in the outer or valence shell.

Atomic Radius Trend

Bigger down a group and from right to left on a period

Ionic Radius trend

Bigger down a group and from right to left on a period

Except for Non-metals (Bigger) and Noble Gasses (Smaller)

Are Cations bigger or smaller than the parent atom?

Smaller, because they have less electrons, less shells, higher attraction between nucleus and outermost electron

Are Anions bigger or smaller than the parent atom?

Larger, because they have more electrons, maybe more shells, lower attraction between nucleus and outermost electron

Electronegativity

The tendency for an atom to attract electrons toward itself

The three most important electronegative elements are Fl, O, and N.

Electronegativity trend

Increases up a group and from left to right on a period.

Effective nuclear charge

Distance of the electrons. 

(=no of protons and electron shells)

Electronegativity difference (define and remember symbol)

The difference in electronegativity between two bonding atoms. It can tell us a rough measure of polarity and thus the bond type.

Classify the elements into metals, nonmetals and metalloids

Metals Properties

1. Shiny

2. Good conductors of heat

3. good conductor of electricity

4. Metal oxides tend to be basic

5. Metallic properties increase down a group

Non-metals properties

1. Non-metal oxides tend to be acidic

2. Important Non-metal oxides

   • Sulfur:

     • Sulfur trioxide, SO₃, reacts with water to form sulfuric acid

     • Sulfur dioxide, S02, and oxides of nitrogen, such as N02, are the main gases responsible for the formation of acid rain

   • Carbon

     • Carbon dioxide, CO2, is acidic. Increasing atmospheric levels are responsible for ocean acidification, which affects marine life.

Ionization energy

Minimum energy needed to remove one mol of electrons from one mol of gaseous atoms. 

Usually an endothermic process

Measured in kJ mol^-1

Ionization Energy Trend

Generally, the values increase across a period because the extra electrons fill the same energy level and the extra protons in the nucleus attract this energy level more closely, making it harder to remove an electron. The values do not increase regularly across a period because new sub-levels are being filled . The p sub-level is higher in energy than the s sub-level. This explains why the value for B (ls22s22p1) is slightly lower than the value for Be (l s22s2) and the value for Al ([Ne]3s23p1) is slightly lower than Mg ([Ne]3s2). There is also a drop in value between N (ls22s22p3) and O (ls22s22p4) and between P ([Ne]3s23p3) and S ([Ne]3s23p4). This is because when electrons pair up in an orbital there is increased repulsion, so the paired electron is easier to remove compared with when the three electrons are all unpaired, one each in the three separate p orbitals.

Electron affinity

Minimum energy released when one mole electrons is added to one mole of gaseous atoms. 

Usually an exothermic process

When the numerical value increases, electron affinity decreases

Electron Affinity trend

Aluminium Special case

Al is a metal but its oxide is amphoteric, that is, it can be either basic or acidic depending on whether it is reacting with an acid or a base

Oxidation state

The number assigned to that atom to show the number of electrons transferred in forming a bond

Tool to identify which species has been oxidized and which reduced in a redox reaction

Help determine the correct formula and name for ionic compounds.

Oxidation state rules:

1. Ionic compound: OS of each element is equal to the charge carried by the ion

2. Covalent compounds: Assume that the compound is ionic, with the more electronegative element forming the negative ion

3. ∑OS’s = Zero

4. ∑OS’s ion = Charge of Ion

5. Elements not combined with other elements have an oxidation state of zero

Group 1 metals Oxidation state

Always +1

Group 2 metals Oxidation State

Always +2

Hydrogen Oxidation state

Usually +1, Metal Hydrides -1

Oxygen Oxidation state

Usually -2, Peroxides (-1) and F₂O

Fluorine Oxidation State

Always -1

Chlorine Oxidation State

Usually -1, Compounds with O or F

Naming of ionic compounds

Metals in groups l, 2, and 13 form cations with oxidation states of +1, +2, and +3, respectively

Transition metals (and other metals, like tin and lead) form more than one ion as they have variable oxidation states. Roman numerals are used for oxidation numbers.

Oxyanions

Formed when oxygen combines with another element to form a complex ion

Should nowadays always have suffix “-ate” and then roman numeral

Complex Ion

Charged species consisting of a central metal atom or ion bonded to one or more molecules or anions, known as ligands. These ligands can be neutral molecules or anions that donate pairs of electrons to the metal, facilitating coordination bonds.