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Download CVB101 – Lecture 3 Chemical Bonding • Chemical bonding
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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 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 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