Elements of Life

Structure of Atom

Structure of Atom

• Positively charged nucleus containing protons and neutrons with negatively charged electrons orbiting around it, in shells.

• Electrons make up the most volume of atoms, atom is mostly empty space.

Relative mass of electrons

Negligible

Atomic number

Number of protons in the nucleus of an atom

Why do all atoms of the same element have same atomic number?

Same number of protons in nucleus

Mass number

Number of protons and neutrons in the nucleus of an atom

Ion

Has different numbers of protons and electrons (protons = electrons)

1- Dalton’s model

Early 19th century - solid spheres and the smallest particles of an element. Elements made from different spheres.

2- J J Thompson

1897 - Plum Pudding model.

• Discovery of electrons and its mass + charge.

• Positively charged sphere with negatively charged electrons embedded into it.

3- Rutherfold

1909 - Nuclear model.

• Discovered nucleus

• Gold leaf experiment - Fired alpha particles at thin sheet of gold. Expected most to be deflected very slightly but most past straight through + some deflected back.

• Positive nucleus in mainly empty space with cloud of electrons

4- Chadwick

Discovered neutrons.

• Fired alpha particles at a thin sheet of Beryllium

5- Bohr

Electrons in fixed energy orbitals.

• When an electron moves between shells, EMR is emitted or absorbed.

• Each shell can only hold a fixed number of electrons.

6- Quantum model

Electrons doesn’t have same energy in shells.

• Contain subshells.

Isotopes

atoms that have the same number of protons but a different number of neutrons.

• Have the same chemical properties as they have the same electron configurations.

• Have slightly different physical properties as they depend on the mass of an atom

Relative atomic mass (A_r)

The average mass of one atom of the element, taking into account all the isotopes and their abundances, relative to carbon-12 = 12g.

Relative isotopic mass

average mass of an isotopic atom of an element, relative to 12g of carbon-12.

Relative molecular mass (M_r)

The average mass of a molecule or formula unit, relative to carbon-12 = 12g

One mole of a substance

contains the same number of particles as there are atoms in 12g of carbon-12.

= is the M_r in grams.

= 6.02 x10²³ (Avogadro constant, N_A)

Empirical formula

The smallest number ratio of atoms in a molecule.

Molecular formula

The actual number of atoms of each element in a molecule

Hydrated solids

• Compounds containing water molecules incorporated into them; water of crystallisation

% yield formula

% yield=  (actual moles of product)/(theoretical moles of product)  ×100

Why % yield isn’t 100%

• Loss from unreacted chemical reagents

• Mass lost in transferring materials with equipment

• Reversible reaction so products turn back into reactants

• Mass lost if gas products escape into surroundings

Electronic structure

Electrons exist in shells which contain subshells - s,p,d,f

• Each subshell consists of different number of atomic orbitals – s=1, p=3, d=5, f=7.

Atomic orbital

A region of space in an atom where there is a high probability of finding the electron.

• Each atomic orbital can hold up to 2 electrons – have opposite spin.

Shells and subshells

1 = 1s = 2e

2 = 2s, 2p = 8e

3 = 3s, 3p, 3d = 18e

4 = 4s, 4p, 4d, 4f = 32e

Orbitals

• s-orbital is spherical-shaped

• p-orbital is dumb-bell shaped

• The closer the average distance of an electron to the nucleus, the lower its energy

• An electron will occupy the lowest energy subshell available

Ionic bonding

• Forms ions when electrons are transferred between atoms

• Strong electrostatic attraction between positively and negatively charged ion.

• Forms giant ionic lattice

Physical properties of ionic compounds

• Solids at room temperature with a high melting point – very strong electrostatic attractions between opposite ions in the lattice need a lot of thermal energy to break.

• Most are water soluble – water molecules are polar which form attractions to the charged ions.

• Conduct electricity when molten or in solution – ions are free to move and carry charge.

Covalent bond

Shared pair of electrons.

Dative bond

Covalent bond in which both electrons came from the same atom

Giant covalent structures vs simple molecules

have stronger electrostatic attraction than simple covalent molecules

Physical properties of simple covalent molecules

• low melting + bp, weak intermolecular forces so little thermal energy required

• Can’t conduct electricity - no overall charge

• Most insoluble in water

Physical properties of giant covalent lattices

• High melting point - strong intermolecular forces in covalent bonds

• Insoluble in any solvent

• Can’t conduct electricity except graphite

• Extremely hard from strong bonds

• Good thermal conductors - vibrations travel easily through lattice

Metallic structures

Positively charged metal ions lattice surrounded by a sea of delocalised electrons with strong attraction.

Physical properties of metals

• High melting point + solid at room temperature - strong electrostatic forces

• Group 2 higher mp than group 1 as ions has higher charge so stronger electrostatic attractions

• Conducts electricity - delocalised electrons are free to move and carry charge

• Ductile, metal ions slide over each other as no bonds holding them together

• Insoluble

Shape of molecules

• Sets of electron pairs around a central atom repelling to get as far apart as possible.

• Lone pairs repel more than bonding pairs

Elements in the same period have

the same number of electron shells occupied

Elements in the same group have

the same number of electrons in their outer shell.

similar physical and chemical properties

Mp of metals increase across the period because

stronger metallic bonds as metal ions have increasing number of delocalised electrons and smaller ionic radius (pulled closer from more protons) so higher charge density thus stronger electrostatic attractions.