CVB101 – Lecture 3 Chemical Bonding • Chemical bonding describes ...

CVB101 – Lecture 3 Chemical Bonding  Chemical bonding describes the exchange and sharing of electrons Chemical Formulae  Chemical symbol – one or two letters which identify each element  Chemical formula – expression of the composition of a compound in terms of chemical symbols  Chemical formulae indicate which elements are present and in what ratio Molecules  Molecule – an aggregate of at least two atoms in a definite arrangement help together by covalent bond  A molecule can contain atoms of the same element or atoms of two or more elements which are in a fixed ratio  law of definite proportions  Polyatomic molecules contain more than two atoms Empirical formula  Empirical formula – an expression with the smallest whole numbers giving the correct ratios of the elements present Structural formula  A representation showing how atoms are bound in a compound Electron arrangements within atoms  Number of protons in the nucleus determines the chemical identity of the atom  Chemical properties, most importantly, chemical reactivity is determined by the electrons, more precisely, electronic structure (number of eincluding their distribution around nucleus and their energies) – explained by quantum theory Electron subshells  S subshell has 2 electrons  P subshell has 6 electrons  D subshell has 10 electrons  F subshell has 14 electrons  Valence electrons: the electrons in the outmost shell (valence shell) of an atom  Core electrons: the electrons in the inner most shells Covalent bonds  Atoms are held together by sharing electrons  This type of bonding typically occurs between atoms of non-metals Ionic bonds

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Ionic bonds are a result of electron transfer between atoms to form ions – electrostatic attraction of positive and negative ions This type of bonding occurs between ionic compounds Ionic bonds are present in compounds of metals and non-metals

Metallic bonds  Metallic bonds occur between metal atoms in metallic solids  The atoms can be though of as held together in a “sea of electrons” Ions    

An atom or group of atoms with a net positive or negative net charge Electrons may be gained or lost Anions – negative charged ion Cations – positive charged ion

Ionic compound  A combination of cations and anions with no overall net charge

MEMORIZE THESE CVB101 – Lecture 4 Absolute and Relative uncertainty  Absolute uncertainty – uncertainty in the measurement expressed in the same unit  Relative uncertainty – compares the size of the uncertainty with the measurement Stoichiometry Blackman chapter 3  Molecular mass – the mass of a compound in atomic mass units (u)  Molecular mass = sum of atomic masses of each atom in its chemical formula  Avogadro’s number: 6.022x1023 mol-1  Molar mass – the mass of 1 mole of something  Molar mass is measured in grams per mole (g mol-1)  How many molecules in 2 mol of water? o Number of atoms = number of moles x avogadro’s number o N=n x NA

CVB101 – Lecture 5 Percent composition  Percent by mass of each element in a compound  Mass % = mass of element in compound/mass of compound x 100  We can calculate theoretical percentage composition from a chemical formula Chemical reactions  A process in which a substance (or substances) is changed into one or more new substances  Involve reorganisation of the atoms in one or more substances  Burning or combustion  Atoms are neither created or destroyed in chemical reactions o All atoms in the reactant must be accounted for in the products o Chemical equations must be balanced

CVB101 – Lecture 6 Reaction Types  Precipitation Reactions  Acid base reactions  Oxidation and reduction reactions Precipitation Reactions  Definition of precipitate – an insoluble solid that separates from a solution  Definition of precipitation reactions – reaction that result in a formation of a precipitate  Include state symbols in chemical equations!! Solubility  The maximum amount of solute that will dissolve in a given quantity of solvent (at a specific temperature)  Some compounds are very soluble e.g. NaCl o It is possible to make very concentrated solutions on NaCl  Other compounds are not very soluble e.g. AgCl o If AgCl solid is placed in water, only very small amount will dissolve. The rest stays as a solid  Saturated solution – a solution in which no more solute will be dissolved Ionic Equation  An equation that shows dissolved species as free ions  Net ionic equations – cancelling spectator ions from both sides gives the net ion equation for the reaction

Acids and bases  Acids: sour taste (acetic acid in vinegar, citric acid in lemons) o Turn litmus from blue to red o Acid solutions conduct electricity  Bases: bitter tastes o Turn litmus from red to blue o Feels slippery (soaps are formed from bases plus fats) o Base solutions conduct electricity Bronsted acids and bases  Arrhenius definition: o Acid – a substance that ionize in water to produce H+ ions o Base – a substance that ionize in water to produce OH- ions  Bronsted definition o Acid – a proton donor o Base – a proton acceptor o Acids and bases do not have to be in aqueous solution

Diprotic and Triprotic acids  Diprotic acids can release two protons  Triprotic acids can release three protons Neutralisation Reactions  Neutralisation reaction – a reaction between an acid and a base  Salt – an ionic compound made up of cation(s) other than H+ and anion(s) other than OHpH 

A measure of the acidity of a solution (i.e. hydrogen ion concentration)

Reduction and Oxidation (redox) reactions  Definition of redox reaction – a reaction in which electrons are transferred  Half reactions are used to show electron transfer.  Oxidation – loss of electrons  Reduction – gain of electrons Common redox reactions  Combination reactions – a reaction in which 2 or more substances combine to form a single product  Decomposition reactions – the breakdown of a compound into 2 or more products

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Combustion reactions – reactions of a substance with O2 usually with release of heat and light

Oxidation numbers  Oxidation number (oxidation state) – the number of charges an atom would have in a molecule (or ionic compound) if electrons were transferred completely  Increased oxidation number, the atom has been oxidised  Decrease oxidation number, the atom has been reduced  The sum of the oxidation numbers for all atoms in a compound is always equal to the total charge of that compound