Y11 Chemistry - Periodic Table Trends and Electronic Configuration

The Periodic Table

An arrangement of the elements to show the patterns of their properties.

Who designed the periodic table?

Dimitri Mendeleev

Period

A row (horizontal) on the periodic table.

Group

A column (vertical) on the periodic table.

Where are nonmetals and metals found on the periodic table?

Nonmetals are generally found to the right and up, whereas Metals are found as you go down and to the right.

Trends in structure across a group

Metallic Network (Na, Mg, Al) → Covalent Network (C, Si) → Discrete molecule (P, S, Cl)

Halogens

Group 17. Very reactive nonmetals that form 1- anions.

Noble Gases

Group 18. Gases that are unreactive/inert with few exceptions. They have 8 valence electrons with the exception of Helium.

Alkali Metals

Group 1 (excl. Hydrogen). Soft, low melting point metals that are highly reactive and react vigorously with water and acids, producing hydrogen gas, they form 1+ cations.

Alkali Earth Metals

Group 2. Only slightly less reactive than their group 1 cousins, they react strongly with acids to produce hydrogen gas. With the exception of Beryllium, they react with water to form a metal hydroxide and hydrogen gas, and from 2+ cations.

Transition Metals

Groups 3-12, also known as d-block.

How do Electrons work in Bohr's model of the atom?

Electrons are found in shells/quantised energy levels that are numbered according to their distance from the nucleus (farther = higher energy)

Maximum number of electrons that can be held in a shell

2(n^2)

Valency

The distance between an element and the nearest noble gas in terms atomic number.

Valence electrons

Electrons in the outermost shell of an atom, they are involved in chemical bonding between atoms

Core/Inner electrons

All electrons that are not in the outermost shell and are therefore not involved in bonding. They shield the valence electrons from nuclear attraction.

Lewis Structure

A diagram of an atom with its atomic symbol surrounded by dots representing valence electrons.

How do ions form?

Atoms either gain (anions) or lose (cations) electrons until they reach the stable electron configuration of the nearest noble gas.

The octet rule

The tendency of atoms to gain, lose, or share electrons until they have 8 (and octet of) electrons in their outermost/valence shell.

What facts can be inferred from Coulomb's law? [F=(kq1q2)/(r^2)]

1. The strength of an electrostatic attraction is directly proportional to the magnitude of the charges involved.

2. The strength of an electrostatic attraction in inversely proportional to the square of the distance between the charges, thus a change to the distance between charges will have a greater effect on the force than an changes of a similar magnitude to size of the charges.

Core Charge/Nuclear Attraction/Nuclear Charge

The force of attraction on the valence electrons by the atom's nucleus.

Core Charge Calculation

No. of protons - no. of inner (non-valence) electrons

Atomic radius

The measure of the size of an atom.

Core Charge and atomic radius trends down a group

Core charge is constant, distance between nucleus and valence electrons increases, atomic radius increases

Core Charge and atomic radius trends across a period (left to right)

Core charge increases, distance between nucleus and valence electrons remains constant, atomic radius decreases

Ionisation

The process of removing a valence electron from an atom to make a positively charged ion.

Ionisation energy

The amount of energy required to remove an electron from an atom in gaseous form.

First Ionisation Energy

The amount energy required to remove the first valence electron from an atom. The more strongly the valence electron is attracted to the nucleus, the more energy is needed to remove it. It increases with an increase in core charge and decreases with an increase in atomic radius.

Factors that effect first Ionisation energy

Nuclear Charge: Greater attraction to the nucleus makes electrons harder to remove.

Atomic Radius: Electrons that are farther from the nucleus are easier to remove.

Inner Electron Shielding: Greater Shielding leads to easier removal of valence electrons

Successive Ionisation energy

The energy required to remove electrons from an atom after the first. Each electron is harder to remove than the last, with a jump in required energy for each shell inwards. Removing more electrons will create an increasingly positively charged ion.

Why are group 2 metals less reactive than group 1?

They need to lose more valence electrons (2 vs 1) to form ions, which requires more energy.

Metals

Elements that are generally, shiny, malleable, ductile, solid at room temperature, and can conduct electricity and heat well

Nonmetals

Elements that usually are gases at room temperature and can not conduct electricity. They tend to be reduced (gain electrons) in chemical reactions.

Metalloids

Elements with both metallic and nonmetallic properties e.g. Silicon.

Metallic character

How close an element is to typical metal properties, such as the ability to lose an electron easily to form a cation.

What is a good reducing agent?

A metallic element that loses an electron and become oxidised when reacting with substances such as acids, water, or oxygen.

Electronegativity

The ability of an atom to attract a bonding pair of electrons to itself in a covalent bond. Typically, nonmetals are highly electronegative, while metals have low electronegativity.

The most and least electronegative elements

Fluorine (4.0) is the most electronegative

Caesium (0.7) is the least electronegative

How does electronegativity relate to reactivity?

For nonmetals, as they want to gain electrons, higher electronegativity makes them more reactive, while those that cannot attract electrons as strongly are less reactive.

However, as metals want to lose electrons, lower electronegativity makes them more reactive, and more electronegative metals will be less reactive.

Ionisation Energy and Electronegativity trends down a group

Ionisation energy and electronegativity decrease, as more shells means that valence electrons are farther away (higher atomic radius) and therefore less attracted to the nucleus.

Ionisation Energy and Electronegativity trends across a period (L to R)

Ionisation energy and electronegativity increase, as more protons means more core charge, leading to the valence electrons being more attracted to the nucleus.

How is covalent bonding capacity influenced by group?

Group 14 elements form 4 covalent bonds, e.g. Carbon.

Group 15 elements form 3, e.g. Nitrogen.

Group 16 elements form 2, e.g. Oxygen.

Group 17 elements form 1, e.g. Fluorine.