2 The Periodic TableDeveloped independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870’s).Didn’t know much about atom.Put in columns by similar properties.Predicted properties of missing elements.
3 DetailsValence electrons- the electrons in the outermost energy levels (not d).Core electrons- the inner electrons
4 7.2 Effective Nuclear Charge Properties of atoms depend not only on their electron configurations but also on how strongly their outer electrons are attracted to the nucleus.We view the net electric field of the nucleus as if it results from a single positive charge. This is known as the EFFECTIVE NUCLEAR CHARGE, Zeff.
5 ShieldingValence electrons are attracted to the nucleus of the atom and are repelled by the other electrons in the atoms.The inner electrons partially shield or screen the outer electrons from the attraction of the nucleus.
6 Periodic Trend of ZeffThe effective nuclear charge increases as we move across any row (period) of the table.The number of core electrons stay the same as we move across the row, the actual nuclear charge increases.
7 Periodic Trend of ZeffGoing down a column, the effective nuclear charge experienced by valence electrons changes far less than it does across a row.We would expect the effective nuclear charge for the outer electrons in lithium and sodium to be about the same, roughly 3-2 = +1 for Li and = +1 for Na.
8 7.3 Sizes of Atoms and IonsQuantum Mechanical Model – no defined boundaries at which the electron distribution becomes zero.We can define atomic size based on the distance between atoms in various situations.Nonbonding atomic radiusBonding atomic radius (shorter)Allows us to estimate the bond lengths between different elements in molecules
9 Practice Use page 231, Figure 7.5: Natural gas used in home heating and cooking is odorless. Because natural gas leaks pose the danger of explosion or suffocation, various smelly substances are added to the gas to allow detection of a leak. One such substance is methyl mercaptan, CH3SH. Predict the lengths of the C-S , C-H, and S-H bonds in this molecule.
10 PracticeUse figure 7.5 to predict which will be greater, the P-Br bond length in PBr3 or the As-Cl bond length in AsCl3.
11 Periodic Trends in Atomic Radii 1. Within each column, atomic radius tends to increase from top to bottom.Results primarily from the increase in the principal quantum number of the outer electrons.2. Within each row, atomic radius tends to decrease from left to right.The increase in the effective nuclear charge as we move across a row steadily draws the valence electrons closer to the nucleus.
12 You Try It… Arrange the following atoms in order of increasing size: P, S, As, SeArrange the following atoms in order of increasing atomic radius:Na, Be, Mg
13 Periodic Trends in Ionic Radii 1. Cations are smaller than their parent atoms2. Anions are larger than their parent atoms.For ions carrying the same charge, size increases as we go down a column in the periodic table.
14 You Try It… Arrange these atoms and ions in order of decreasing size: Mg2+, Ca2+, CaWhich of the following atoms and ions is the largest?S-2, S, O-2
15 Isoelectronic seriesA group of ions all containing the same number of electronsO-2, F-, Na+, Mg2+, Al3+ all have 10 electrons.We can list the members in order of increasing atomic number, and therefore nuclear charge increases as we move through the series. Because the number of electrons is constant, the radius of the ion decreases with increasing nuclear charge, as the electrons are more strongly attracted to the nucleus.Oxide is the largest ion, smallest atomic numberaluminum the smallest ion, highest atomic number
16 You Try It…Arrange the ions of potassium, chloride, calcium, and sulfide in order of decreasing size.
17 7.4 Ionization EnergyIonization energy-the energy required to remove an electron from a gaseous atomHighest energy electron removed first.First ionization energy (I1) is that required to remove the first electron.Second ionization energy (I2) - the second electronetc. etc.
18 Trends in ionization energy for MgI1 = 735 kJ/moleI2 = 1445 kJ/moleI3 = 7730 kJ/moleThe effective nuclear charge increases as you remove electrons.It takes much more energy to remove a core electron than a valence electron because there is less shielding.
19 Explain this trend For Al I1 = 580 kJ/mole I2 = 1815 kJ/mole
20 Across a Period Generally from left to right, I1 increases because there is a greater nuclear charge with the same shielding.As you go down a group I1 decreases because electrons are farther away.
21 It is not that simpleZeff changes as you go across a period, so will I1Half filled and filled orbitals are harder to remove electrons from.here’s what it looks like.
25 Periodic Trends in 1st Ionization Energies 1. Within each row, I1 generally increases with increasing atomic number.Alkali metals show the lowest, noble gases the highest.2. Within each column, the ionization energy generally decreases with increasing atomic number.He>Ne>Ar>Kr>Xe
26 You Try It…Referring to a periodic table, arrange the following atoms in order of increasing first ionization energyNe, Na, P, Ar, KWrite the electron configuration of:Calcium IonCobalt (III) IonSulfide Ion
27 7.5 Electron AffinitiesThe energy change that occurs when an electron is added to a gaseous atomAffinity = attractionCl(g) Cl+ + e- Delta E = 1251(first ionization energy) kJ/molCl(g) + e- Cl- Delta E = -349 kJ/mol(electron affinity)
28 RememberIonization energy measures the ease with which an atom LOSES an electron.Electron affinity measures the ease with which an atom GAINS an electron.
29 Electron Affinity - the energy change associated with the addition of an electron Affinity tends to increase across a periodAffinity tends to decrease as you go downin a periodElectrons farther from the nucleusexperience less nuclear attractionSome irregularities due to repulsiveforces in the relatively small p orbitalsValues of Electron affinity are always negative, the more negative the value the more stable the atom becomes. Therefore, the more negative the value, the more that atom wants an electron.
31 Electronegativity A measure of the ability of an atom in a chemical compound to attract electronsElectronegativities tend to increase acrossa periodElectronegativities tend to decrease down agroup or remain the same