Atomic Structure, Bonding and Periodicity Contents • Atomic Structure • Amount of Substance • Bonding • Periodicity Atomic Structure • Fundamental Particles • Mass Number and Isotopes • Mass Spectrometer • Electron Arrangement • Trends Ionisation Energies Fundamental Particles Atoms have a small central nucleus made up of protons and neutrons around which there are electrons. Neutron Electron Atomic Relative Relative Particle Mass Charge + neutron 1 0 + proton 1 +1 electron negligible -1 Proton Isotopes • Atoms of the same Mass number: this is the number of element always have the protons and neutrons (A) same number of protons. • Isotopes are atoms of the same element but with 4 different number of neutrons. This gives rise to different mass numbers. 2 He Helium • Relative abundance is the amount of each isotope as Atomic number: this is the the percentage for that number of protons (Z) element occurring on the Earth Mass Spectrometer A curved magnet An Charged The deflects the electron plates The ions are detected sample is positive ions. The gun accelerate to produce a mass put in via lighter ions are ionises positive spectrum an injector deflected more the atoms ions Ionisation acceleration deflection detection •The mass spectrum gives the relative Example: Copper masses of each isotope and the 69% of copper atoms have a abundance of each isotope. relative mass of 63. •The relative atomic mass is the 31% of copper atoms have a weighted average mass of an atom of an relative mass of 65. element compared with 1/12 of the mass The weighted average is calculated of 12C. as follows: •We can calculate this from mass (0.69 x 63) + (0.31 x 65) =63.62 spectra Arrangement of Electrons • Energy Levels or Shells – The simplest model of electrons has them orbiting in shells around the nucleus. Each successive shell is further from the nucleus and has a greater energy. • Sub Shells and Orbitals – This model can be further refined by the concept of sub shells and orbitals. – Sub shells are known by letters s, p, d, and f. The s sub shell can contain 2 electrons, p 6, d 10 and f 14. – Electrons occupy negative charge clouds called orbitals, each orbital can hold only 2 electrons. Each type of shell has a different type of orbital. Arrangement of Electrons • How we write electron configurations – Electrons fill the lowest energy level first this means it is generally easy to predict how the electrons will fill the orbitals (it gets more complicated with the transition metals). Element Electron Element Electron configuration configuration H 1s Ne 1s22s22p6 He 1s2 Na 1s22s22p63s1 Li 1s22s Mg 1s22s22p63s2 Be 1s22s2 Al 1s22s22p63s23p3 B 1s22s22p1 Si 1s22s22p63s23p3 C 1s22s22p2 P 1s22s22p63s23p3 N 1s22s22p3 S 1s22s22p63s23p4 O 1s22s22p4 Cl 1s22s22p63s23p5 F 1s22s22p5 Ar 1s22s22p63s23p6 Trends in Ionisation Energies The first ionisation energies of group 2 First ionisation energies of period 3 electrons elements 1000 2000 Energy (Kj/mol) Energy (Kj/mol) 800 1500 600 1000 400 200 500 0 0 Be Mg Ca Sr Ba Na Mg Al Si P S Cl Ar Group 2 elements Period 3 elements Going down a group in the periodic table there are more filled energy levels between the Going across a period of elements: nucleus and the outer most electrons there are more protons in each nucleus so the these shield the outer electrons from the nuclear charge in each element increases, attractive force of the positive nucleus. this increases the attractive force acting on as the radius of the atom increases, the distance the outer most electrons. between the nucleus and the outer electron there is no significant increase in shielding as increases and therefore the force of attraction each successive electron enters the same between the nucleus and outer most electrons is energy level as the one before. reduced. Overall more energy is needed to remove the These factors mean that less energy is needed first electron from its outermost shell. to remove the first electron from an atom at the bottom of the group compared to one at the top of the group. Amount of Substance • Amount of Substance • Calculations • Balancing Equations • Reacting Quantities Amount of Substance • Different atoms have different masses. 1g of carbon has far fewer atoms than 1g of hydrogen atoms. Chemists need a method of quantifying atoms. • We use a quantity called the amount of substance which is measured in moles. • One mole contains 6.02 x 1023 particles. • The relative atomic mass is the mass of one mole of that element’s atoms. • The relative molecular mass is the mass of one mole of molecules Calculations • Amount of substance, n • The Ideal Gas equation – n = mass Mr • Providing that the pressure and temperature • Solution calculations of gases are the same, – The concentration of equal volumes of two solutions are measured in gases can be assumed to mol dm3 • c= concentration in dm3 have the same number of • V = volume in cm3 moles. – n = Vc – pV = nRT 1000 • p is the pressure in Pa • V is the volume in m3 • T is the temperature in P Kelvin r • R= 8.31 JK-1 mol-1 Balancing Equations • Chemical reactions involve the rearrangement of atoms not the making or destroying of atoms. • It is necessary to make sure that you have the same amount of atoms on both sides of the equation. • State symbols can also be added to show the physical condition of the reactants and products • (s) – solid, (l) – liquid, (g) – gas, (aq) – aqueous Reacting Quantities • The numbers in a balanced equation give the ratio of the amount of each substance in the reaction. We can use this information to calculate quantities of reactants or products. • 50g of CaCO3 are heated how much CaO will be formed • First write a balanced equation: – CaCO3(s) CaO(s) + CO2(g) • Then calculate the Mr of the substances we are interested in: – CaCO3 40 + 12 + (3 x 16) = 100 – CaO 40 + 16 =56 • Calculate the number of moles of CaCO3 used. – n=mass/Mr 50/100 = 0.5 mol • From the equation we can see that one mole of CaCO3 produces one mole of CaO. Therefore 0.5 mol of CaCO3 produces 0.5 mol of CaO. • Finally calculate the mass of 0.5 mol of CaO: – n=mass/Mr, mass=Mr x n, 0.5 x 56 = 28g • Therefore 50g of CaCO3 produces 28g CaO Bonding • The nature of bonds • Bond polarity and the polarisation of ions • Intermolecular forces • Hydrogen Bonding The Nature of Bonds Covalent Bonding H • When non-metals react together both atoms need to gain electrons to obtain a full shell of electrons they do this by forming a covalent bond. H C H • The atoms are held together by shared pairs of electrons from the highest energy level of both atoms. H Ionic bonding Atoms lose or gain electrons to attain a complete outer + _ shell of electrons. An ionic bond is formed when electrons are lost and gained by two or more atoms. Na Cl When atoms lose electrons they become positive ions, when they gain electrons they become negative ions. It is the electrostatic forces of attraction which hold the ions together Metallic Bonding In metals, positive metal ions are held - - - +- - + -+ + - - ++ - + - + together by electron clouds. These + - + + - +- electrons-are+ to-move through the free + - + + - -+ - + +- + + this is why metals conduct + - - + + + - structure, - + + - + electricity. Bond Polarity and the Polarisation of Ions • In reality not all bonds are perfectly covalent or ionic. To explain why we have to define a concept called electronegativity. • Electronegativity is the ability of an atom to attract the bonding electrons. • In hydrogen fluoride the fluorine atoms are much more electronegative than the hydrogen. It pulls the electrons toward it creating what is called a polar bond. • Ionic bonds can also show polarity, this can happen if the electron cloud is distorted by strong charges on one of the ions. If a cation is highly charged it will exert a strong electrostatic attraction on the anion and distort the electron cloud. If the anion has a large electron cloud it will be easily distorted. If the electron cloud is distorted there will be electron density between the two ions giving the bond some covalent characteristics • Molecules with asymmetric charge distribution are said to be polar molecules and to possess a dipole Intermolecular Forces • Permanent dipole- permanent dipole interactions occur d+ H Cl d- between polar molecules. This happens when the negative end of one Weak electrostatic forces molecule is attracted to the positive end of another. This force is d- Cl H d+ much weaker than intramolecular bonding •Temporary dipole –induced dipole interactions exist between non-polar molecules and monatomic species such as the noble gases. The distribution of the electron cloud on a molecule is not constant and at any given time it can asymmetric. this confers a temporary asymmetry on the charge distribution. The molecule is said to possess a temporary dipole, this temporary dipole can induce another temporary dipole in an adjacent molecule. There is a resulting weak electrostatic force between the two molecules d- Hydrogen Bonding O • Hydrogen bonds are a special case of H d+ H d+ permanent dipole-permanent dipole d- bonding. O • It exists where an electronegative element such as oxygen, chlorine H d+ H d+ fluorine or nitrogen is bonded to d- hydrogen. O • Hydrogen bonding causes stronger intermolecular bonds than would H d+ H d+ otherwise be predicted this increases d- the boiling point of substances such as O water. H d+ H d+ Periodicity • Chemists classify elements according to their position in the periodic table. • Periodicity is the term used to describe the repeating pattern of properties observed within the periodic table. H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac s- block d-block p-block Trends in Group 2 Compounds • Progressing down group 2 the atomic radius increases due to the extra shell of electrons for each element. • Going down the group the first ionisation energy decreases there is more shielding between the nucleus and the outer electrons and the distance between the nucleus and the outer electron increases and therefore the force of attraction between the nucleus and outer most electrons is reduced. • Generally the melting point of the metals decreases down the group this is because as the metal ions get larger the distance between the bonding electrons and the positive nucleus gets larger and reduces the overall attraction between the two. For similar reasons the electronegativity decreases. • The reactions of the elements with water become more vigorous down the group. When they do react they produce hydroxides and hydrogen. • The solubilities of the hydroxides of the elements increase going down the group. • The solubilities of the sulphates of the elements decreases down the group. • Barium sulphate is insoluble and is used as a qualitative test to identify sulphate ions. Trends in Period Three of the Periodic Table Property Trend from Explanation left to right Atomic radius decreases because the nuclear charge increases First ionisation increases because the nuclear charge energy increases Electronegativity increases because the nuclear charge increases Electrical Increases until because the metals have an conductivity the non metals increased number of delocalised electrons Boiling point and Increases until Because these properties depend the middle then on the forces between the particles. Melting point decreases This depends on the structure of the element which varies from metallic to giant covalent to simple molecular. Summary • Atomic Structure – We consider atoms to be formed from three fundamental particles, we can determine the relative atomic mass using a mass spectrometer. • Amount of Substance – Chemists use this concept to count atoms. Using this concept we can calculate reacting quantities, in a given reaction. • Bonding – Bonding within molecules can be described as covalent, ionic or metallic. Often a bond is a hybrid between ionic and covalent. • Periodicity – Trends within the periodic table can frequently be explained by concepts of nuclear charge and shielding.