Fall 2016_CH1010_Dr. Kreider-Mueller CH1010 Exam #2 Study Guide For reference see “Chemistry: An Atoms-focused Approach” by Gilbert, Kirss, and Foster Chapter 3: Atomic Structure, Explaining the Properties of Elements Trends to know (and be able to explain the trend…think about Zeff): o Effective Nuclear Charge (Zeff): the attraction toward the nucleus experienced by an electron in an atom; the positive charge on the nucleus reduced by the extent to which other electron in the atom shield the electron from the nucleus o Ionic radius (Figure 3.35 in textbook) . When you form a cation, the radius shrinks (ex: the ionic radius of Na+ is smaller than the atomic radius of Na). Why? . When you form an anion, the radius expands (ex: the ionic radius of Cl− is larger than the atomic radius of Cl). Why? o Ionization Energy . M + Energy M+ + e− M+ + Energy M2+ + e− M2+ + Energy M3+ + e− …and so forth . Look at Figure Table 3.2 in your book (Think about when/why there are large jumps between successive ionization energies) . Be able to make predictions about ionization energies for various elements o Electron Affinity . Which group in the periodic table has the highest (most negative) electron affinity? Which group(s) in the periodic table have the lowest (values close to zero) electron affinity . Look at Figure 3.37 in your book . Be able to make predictions about electron affinities for various elements . Group 1A elements have low electron affinities so they tend to lose their ns1 electron . Group 2A elements have low electron affinities so they tend to lose their ns2 electrons . Group 7A elements have large electron affinities so they tend to gain one electron, adopting the electron configuration of the neighboring noble gas Group 8A elements are inert, and don’t generally gain or lose electrons Chapter 4: Chemical Bonding Definitions to know: o Molecule: unit of matter that results when 2 or more atoms are joined by covalent bonds. o Ionic Bond: a bond resulting from the electrostatic attraction of a cation for an anion. o Nonpolar Covalent Bond: a bond characterized by an even distribution of charge; electrons in the bonds are shared equally by the two atoms. o Polar Covalent Bond: a bond resulting from unequal sharing of bonding pairs of electrons between two atoms. o Polyatomic ions: charged group of more than one kind of atom joined together by covalent bonds. 1 of 11 Fall 2016_CH1010_Dr. Kreider-Mueller o Electrostatic Potential: the energy a charged particle has due to its position relative to another charged particle. Directly proportional to the product of the charges of the particles . Inversely proportional to the distance between the particles o Lattice Energy: the energy released when 1 mole of an ionic compound forms from its free ions in the gas phase. o Bond Energy: the energy needed to break 1 mole of a particular covalent bond in a molecule or polyatomic ion in the gas phase. o Octet Rule: atoms of main group elements make bonds by gaining, losing, or sharing electrons to achieve a valence shell containing 8 electrons, or four electron pairs. o Bond Length: distance between the nuclear centers of two atoms joined together in a bond. o Bonding Pair: a pair of electrons shared between two atoms. o Lone Pair: a pair of electrons that is not shared between two atoms. o Electronegativity: a relative measure of the ability of an atom in a bond to attract electrons to itself when bonded to another atom. o Resonance: characteristic of electron distributions when two or more equivalent Lewis structures can be drawn for one compound. o Resonance Structure: one of two or more Lewis structures with the same arrangement of atoms but different arrangements of bonding pairs of electrons o Resonance Stabilization: the stability of a molecular structure due to delocalization of its electrons. o Formal Charge: value calculated for an atom in a molecule or polyatomic ion by determining the difference between the # of valence electrons in the free atom & the sum of lone-pair electrons plus half of the electrons in the atom’s bonding pairs. o Bond Order: the number of bonds between atoms: 1 for a single bond, 2 for a double bond, 3 for a triple bond. Be able to Distinguish between the three different types of bonds * Energy is released when bonds are formed, exothermic process* * Energy is absorbed when bonds are broken, endothermic process* 1) Nonpolar Covalent Bond: two atoms evenly share 2 electrons o Length and Strength of covalent bonds: . Double bond is shorter and stronger than a single bond . Triple bond is shorter and stronger than a double bond . The internuclear distance that corresponds to a potential energy minimum (point “c” on the graph below) is the bond length between two atoms. 2 of 11 Fall 2016_CH1010_Dr. Kreider-Mueller 2) Polar Covalent Bond: two atoms share 2 electrons unevenly (unequal sharing). The electrons are more attracted to the more electronegative atom. 3) Ionic Bond: complete transfer of one or more electrons from one atom to another occurs resulting in the formation of 2 charged particles (ions) o Ions of opposite charge attract one another . Metals tend to form cations . We use a Roman numeral in parentheses to indicate # of charges on a cation 2+ Ex: in Fe(II)Cl2 the iron has a 2+ charge, Fe 3+ Ex: in Fe(III)Cl3 the iron has a 3+ charge, Fe . Nonmetals tend to form anions . Common polyatomic ions you should know: + − − NH4 , ammonium NO3 , nitrate ClO4 , perchlorate − − 3− CH3CO2 , acetate OH , hydroxide PO4 , phosphate − − 2− CN , cyanide MnO4 , permanganate SO3 , sulfite − 2− 2− NO2 , nitrite CO3 , carbonate SO4 , sulfate Be able to distinguish between ionic & molecular compounds in molecular level representations Ionic Covalent Be able to account for the differences in physical properties of ionic & covalent compounds. For ex, see table 4.1 from General Chemistry: Atom’s First by McMurry & Fay Trends to know (and be able to explain…think about Zeff): o Lattice Energy . Depends on the charge of the ions and the size (radius) of the ions. The smaller the radius, the higher the lattice energy The larger the charge on the ions, the higher the lattice energy . The higher the lattice energy, the higher the melting point of the ionic solid 3 of 11 Fall 2016_CH1010_Dr. Kreider-Mueller . Be able to predict the magnitudes of lattice energy based on analogous compounds, ionic charge, & ionic radius For example, see the table below: o Electronegativity . Look at Figures 4.8 & 4.9 in your textbook . Be able to predict the ionic/covalent character of bonds based on electronegativity differences Bonds between atoms w/ similar electronegativities are usually nonpolar covalent Bonds between atoms whose electronegativities differ by more than 2 units are largely ionic Bonds between atoms whose electronegativities differ by more than 0.4 units but less than 2 units are usually polar covalent . Fluorine is the most electronegative (value of 4.0) atom . Oxygen is the second most electronegative (value of 3.5) atom Bonding in molecules can be represented by using only valence electrons in electron dot (Lewis) diagrams. Be able to draw electron dot structures: o Section 4.3 of your book gives some steps for drawing electron dot structures o Figure out how many valence electrons you need to account for in your structure o Connect the atoms appropriately o Remember the following: . Hydrogen does not follow the octet rule (only surrounded by 2 electrons) . C, O, N, and F follow the octet rule (boron often does not) . If an element is in period 3 or below, then you can expand the octet if necessary . Carbon likes to form 4 bonds . Nitrogen typically likes to form 3 bonds . Oxygen typically likes to form 2 bonds . Halogens typically like to form 1 bond o Consider if you need any multiple bonds 4 of 11 Fall 2016_CH1010_Dr. Kreider-Mueller o If there is more than one possible structure, you need to be able to determine which structure is best based on the octet rule and formal charges Formal Charge = (# of valence e−) – ½ (# of bonding e−) – (# of nonbonding e−) . The more formal charges there are on the atoms in your structure, the less the structure contributes to the overall structure of the molecule (more formal charge = worse structure) o Consider if you can draw resonance structures for your compound (ex. Ozone, O3) . If you can draw different resonance structures, then the electrons are delocalized within the compound . Delocalization reduces electrons’ potential energy making the molecule more stable . Be able to calculate bond orders for resonance hybrids Chapter 5: Bonding Theories, Explaining Molecular Geometry *Bonding can be described using different models…each is useful & each has drawbacks. * Definitions to know: o Bond Angle: the angle (in degrees) defined by lines joining the centers of two atoms to a third atom to which they are chemically bonded. o Electronic Geometry: the 3D arrangement of bonding pairs & lone pairs of electrons about a central atom. o Molecular Geometry: the 3D arrangement of the atoms in a molecule. o Bond Dipole: separation of electrical charge created when atoms with different electronegativities form a covalent bond. o Hybridization: in valence bond theory, the mixing of atomic orbitals to generate new sets of orbitals that then are available to form covalent bonds with other atoms. o Hybrid Atomic Orbital: in valence bond theory, one of a set of equivalent orbitals about an atom created when specific atomic orbitals are mixed. o Molecular Orbital: a region of characteristic shape and energy where electrons in a molecule are located.