Periodic table and molecules

Flame tests

• The closest shell to the nucleus has the lowest energy level

• When atoms are heated, electrons are given more energy

• They can absorb enough energy that they ‘jump’ up an energy level into the next energy level (shell) 

• When the electron cools back down and returns to its original energy level, it gives off energy in the form of light

• If this light is passed through a prism, the resultant pattern of lines is called an emission spectrum where each line corresponds to light of a different energy level.

• Emission spectra are like fingerprints for elements, meaning that the pattern of lines observed is unique depending on the element.

The Periodic Table

• Elements in the same group react the same way because they have the same number of valence electrons

• A chemical reaction involves either giving, receiving or sharing electrons.

• Alkali metals: the most reactive metals because they have one valence electron

• Alkali earth metals: they are very reactive and have a strong reaction with water

• Transition metals: found in the centre of the periodic table

• Metalloids: staircase between metals and nonmetals

• Non metals: most are gases at room temperature and Bromine is a liquid.

• Halogens: react with metals to form a salt (group 7)

• Noble Gases: unreactive because they have a full valence shell

Trends in the Periodic Table

• Reactivity: the measure of how easily an element wants to give away an electron. Increases down a group and decreases across a period

• Electronegativity: how strongly an atom can attract electrons to its nucleus. Decreases down a group and increases across a period

• Atomic Radius: the size of atoms is measured by their radius. The more protons and electrons an atom has, the larger the electrostatic charge and attraction between them. This causes the atoms to be smaller when there are more protons and electrons because the valence electrons are pulled closer to the nucleus. Increases down a group because there are more electron shells (valence electrons are therefore further away)  and decreases across a period because there are more protons and valence electrons (increase in electrostatic attraction) - opposite of ionisation energy

• Metallic Character: Metals tend to lose electrons more readily because they have fewer in their valence shell. The more readily an element loses valence electrons, the more metallic it is. Increases down a group (valence electrons easier to lose because they are further from the nucleus), decreases across a period (more valence electrons)

• Ionisation energy: the amount of energy required to remove an electron from an element to form an ion. Noble gases have the highest and alkali metals have the lowest. Increases as you travel up and right. (opposite of atomic radius)

Ionic Bonding (metal + non metal)

• Metals give away electrons in order to have  a full valence shell, therefore become positively charged. (cations)

• Non metals gain electrons to have a full valence shell, therefore become negatively charged (anions)

• When metal cations and nonmetal anions attract, it's called an ionic bond.

• Sodium (metal) loses one electron and chlorine (non metal) needs to gain an electron, so sodium gives its electron to chlorine and they become NaCl (Sodium chloride) from Na+ and Cl-

• Ionic compounds are called salts

• Ions form large crystal lattices which are a regular arrangement of (billions of) alternating cations and anions. This results in an electrically neutral ionic salt

• Polyatomic ions: made up of more than one non metal atom and needs to gain or lose an electron in order for the molecule to be stable. The Sulphate ion (SO₄^2-  ) is made from 1 sulphur atom and 4 oxygen atoms chemically bonded together.

Metallic Bonding (metal (cation) + delocalised electrons)

• A metal:  any element that loses its valence electrons to become stable (becomes a cation). Common properties are that they are good conductors of heat and electricity, malleable (hammered, rolled or bent), ductile (drawn into wires), lustrous, have a range of melting points but are mostly high and have high densities

• Metalling bonding: the electrostatic attraction between positive ions and delocalised electrons

• Most metals used in society are used as alloys as combining different metals allows for the selection of particular properties. Metals can also be modified by heat treatments or through the application of coatings

• Alloys: formed by mixing a metal with other metals. They are prepared by melting the metals together and cooling the mixture to form a new solid material. Sometimes nonmetallic atoms such as carbon and silicon are used to form an alloy. Iron + carbon = steel

• Work hardening: hammering or working cold metals which causes a rearrangement of the atoms and a hardening of the metal. The atoms in the metal are altered and smaller grains are formed. The metal becomes more harder but also more brittle

• Heat treatment: the properties of metals can be altered by a process of controlled heating and cooling.

  • Annealing: the metal is heated to a moderate temperature and then is left to cool slowly, allowing larger groups of atoms to form. The metal becomes softer and more ductile.

  • Quenching: the metal is heated to a moderate temperature and then cooled quickly in water, allowing tiny groups of atoms to form. The metal becomes harder but also brittle.

  • Tempering: Quenched metals are then annealed, reducing the brittleness of the material whilst retaining the hardness.

Covalent Bonding

• Covalent bonds are formed when non metal atoms share electrons. Non metals have high electronegativity so they attract electrons easily but don’t give electrons up easily.

• Intramolecular bonds: strong forces of attraction that hold atoms together within molecules

• Intermolecular bonds: weak forces of attraction between molecules (i.e. hold the solid molecular lattice together)

• By sharing the single unpaired electrons, both atoms can stabilise their outer shells

• Covalent molecules tend to be discrete i.e. they don’t form repeating crystal lattice structures, but exist as independent molecules e.g water

• Covalent molecules have unique properties:

  •  Low melting and boiling points (liquids or gases at room temperature) therefore the forces of attraction between particles must be very weak

  • Poor conductors of electricity (no charged particles are free to move through the lattice)

• Bonding electrons: electrons shared between the atoms

• Non-Bonding electrons: the outer shell electrons not involved in forming a bond

• Lone Pairs: pairs of non bonding electrons

• Double and Triple Bonds: Sometimes an element needs to share more than one electron to obtain a full valence shell. If an element shares two electrons with another element it’s a Double bond, if it shares three electrons it’s a Triple bond. Elements cannot share four electrons as the repulsion between the electrons is too great.

Shapes of Molecules

• Valence Shell Electron Pair Repulsion (VSEPR): a model in chemistry which is used to predict the 3D shape of individual molecules. Molecules form 3D shapes because negatively charged valence electrons in the atoms repel each other. VSEPR is based upon minimising the extent of electron pairs repelling each other around a central atom.

• Tetrahedral: 4 bonding pairs. This geometrical arrangement of 4 covalent bonds keeps electron pairs as far apart as possible. 

• Trigonal Pyramidal: 3 bonding pairs and 1 lone pair

• V-Shaped (bent): 2 bonding pairs and 2 lone pairs

• Linear: 1 bonding pair

• Trigonal planar: 3 bonding pairs, no lone pairs (angle of 120 between each bond)