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CHEMICAL BONDING Unit 4 LIGHT AND QUANTIZED ENERGY Nuclear atom and unanswered questions Scientists found Rutherford’s nuclear atomic model fundamentally incomplete Did not explain how electrons are arranged Did not address why negatively charged electrons are not pulled into the atom’s positively charged nucleus Wave nature of light Electromagnetic radiation – form of energy that exhibits wavelike behaviors Visible light Microwaves X-rays Radio/TV waves Wavelength (λ) – shortest distance between equivalent points on a continuous wave Frequency (ν) – number of waves that pass a given point per second Measured in hertz (Hz) Ex: 652 Hz – 652 waves/s Amplitude – wave’s height from the origin to a crest, or from the origin to a trough All electromagnetic waves travel at a speed of 3.00x108 m/s in a vacuum Speed of light is represented by “c” c = λ ν Wavelength and frequency are inversely related Sunlight passing through a prism is separated into a continuous spectrum of colors Particle nature of light The quantum concept – Max Planck concluded that matter can gain or lose energy only in small, specific amounts called quanta Quantum – minimum amount of energy that can be gained or lost by an atom Equantum = hv where “E” is energy, “h” is Planck’s constant, and “v” is velocity Planck’s constant = 6.626x10-34 J·s Planck’s theory: for a given frequency, v, matter can emit or absorb energy only in whole-number multiples of hv (1hv, 2hv, 3hv, etc.) Analogous to child building a wall with wooden blocks The photoelectric effect – electrons (photoelectrons) are emitted from a metal’s surface when light of a certain frequency shines on the surface Albert Einstein proposed light has both wavelike and particlelike characteristics Photon – particle of electromagnetic radiation with no mass that carries a quantum of energy Photon’s energy depends on its frequency Ephoton = hv Atomic emission spectra – set of frequencies of the electromagnetic waves emitted by atoms of the element Each element’s atomic emission spectrum is unique and can be used to determine if that element is part of an unknown compound Neon - light is produced by passing electricity through a tube filled w/ neon gas Neon’s atomic emission spectrum consists of several individual lines of color, not a continuous range of colors as seen in the visible spectrum QUANTUM THEORY AND THE ATOM Bohr model of the atom – Neils Bohr proposed that elements’ atomic emission spectra are discontinuous Energy states of hydrogen Ground state – lowest allowable energy state of an atom Excited state – when an atom gains energy The smaller the electron’s orbit the lower the atom’s energy state Hydrogen’s line spectrum - when in the excited state the electron can drop from the higher-energy orbit to a lower-energy orbit and the atom emits a photon Fig 5-10 Quantum mechanical model of the atom – Louis de Broglie accounted for fixed energy levels of Bohr’s model Electrons behave as waves Only half-wavelengths are possible on a guitar b/c the string is fixed at both ends Only whole numbers of wavelengths are allowed in a circular orbit of fixed radius Fig 5-11 Heisenberg Uncertainty Principle – states it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time Impossible to measure an object w/o disturbing it Tried to measure electrons w/ light but b/c a photon has about the same energy as an electron, the interaction changes the electron’s position Quantum mechanical model of the atom – electrons are treated as waves Atomic orbital - 3-dimensional region around nucleus that describes the electron’s probable location Fig 5-13 Hydrogen’s atomic orbitals Principal quantum numbers – indicate the relative sizes and energies of atomic orbitals As “n” increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases Principal energy levels - atom’s major energy levels (specified as “n”) Energy sublevels Principal energy level 1 consists of a single sublevel; principal energy level 2 consists of 2 sublevels, etc. Sublevels are labeled s, p, d, or f according to the shapes of the atom’s orbitals Each orbital may contain at most 2 electrons All “s” orbitals are spherical All “p” orbitals are dumbbell shaped Not all “d” or “f” orbitals have the same shape Fig 5-15 and 5-16 Table 5-2 ELECTRON CONFIGURATIONS Ground-state electron configurations Electrons tend to assume the arrangement that gives the atom the lowest possible energy The aufbau principle – states that each electron occupies the lowest orbital available All orbitals related to an energy sublevel are of equal energy Energy sublevels within a principle energy level have different energies (Fig 5-17) In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d. and f Orbitals related to energy sublevels within one principal energy level can overlap orbitals related to energy sublevels within another principal level The Pauli exclusion principal – states that a maximum of 2 electrons may occupy a single atomic orbital, but only if the electrons have opposite spins Hund’s rule – states that single electrons w/ the same spin must occupy each equal-energy orbital before additional electrons w/ opposite spins can occupy the same orbitals Orbital diagrams and electron configuration notations Orbital diagram - boxes with zero, one, or two arrows represent orbitals Electron configuration notation – designates the principal energy level and energy sublevel associated w/ each of the atom’s orbitals and includes a superscript representing the number of electrons in the orbital Noble-gas notation Valence electrons – electrons in the atom’s outermost orbitals Electron-dot structure – consists of the element’s symbol, which represents the atomic nucleus and innerlevel electrons, surrounded by dots representing the atom’s valence electrons (Lewis dot structure) FORMING CHEMICAL BONDS Chemical bond – force that holds 2 atoms together Amount of reactivity is directly related to valence electrons Formation of positive ions Cation – positively charged ion Group 1A elements lose 1 valence e-, forming an ion with a 1+ charge Ex: By losing an e-, Na acquires the stable outer electron configuration of Ne Group 2A elements lose 2 valence e-, forming an ion with a 2+ charge Transition metals commonly lose 2 valence e-, forming 2+ ions; however it’s also possible to lose an additional “d” electron, forming 3+ ion Formation of negative ions Anion – negative ion To designate anions, -ide is added to root of element Ex: chloride, sulfide, etc. Nonmetals form a stable outer electron configuration by gaining e Group 5A gain 3 e-, forming ions w/ 3- charge Group 6A gain 2 e-, forming ions w/ 2- charge FORMATION AND NATURE OF IONIC BONDS Formation of an ionic bond Ionic bond – electrostatic force that holds oppositely charged particles together # of e- lost must equal # of e- gained Oxide – ionic bond between metals and oxygen Ex: Ca and F form the ionic compound CaF2 Criss-cross method Properties of ionic compounds Strong attraction of positive ions and negative ions in an ionic compound results in a crystal lattice Solid ionic compounds are nonconductors of electricity b/c of the fixed positions of the ions Liquid ionic compounds (or those dissolved in water) are conductors of electricity b/c ions are free to move Electrolyte – an ionic compound whose aqueous solution conducts an electric current Energy and the ionic bond Endothermic – energy is absorbed during a chemical reaction Exothermic – energy is released during a chemical rxn Formation of ionic compounds from positive and negative ions is always exothermic Attraction of the positive ion for the negative ions close to it forms a more stable system that is lower in energy than the individual ions Lattice energy – energy required to separate one mole of the ions of an ionic compound The more negative the lattice energy, the stronger the force of attraction Directly related to size of the ions bonded Smaller ions have a more negative value b/c nucleus is closer to, and thus has more attraction for, the valence e- Affected by the charge of the ion Ions w/ larger positive or negative charges have a more negative lattice energy NAMES AND FORMULAS FOR IONIC COMPOUNDS Formulas for ionic compounds Formula unit – simplest ratio of the ions represented in an ionic compound Ex: KBr – 1:1 ratio; MgCl2 – 1:2 ratio Determining charge Binary ionic compound – ionic compound formed by 2 ions (metal and nonmetal) Monatomic ion – one-atom ion Ex: Mg2+ and Br Table 8-5 lists common ions of transition metals and groups 3A and 4A Oxidation number – charge of a monatomic ion Compounds that contain polyatomic ions Polyatomic ions – ions made up of more than one atom A polyatomic ion acts as an individual ion Exist as a unit, so NEVER change the subscripts of the atoms w/in the ion Table 8-6 lists common polyatomic ions Naming ions and ionic compounds Most polyatomic ions are oxyanions – polyatomic ion composed of an element, usually a nonmetal, bonded to one or more O atoms Rules for naming nonmetal-oxyanions (ex: N, S): Ion w/ more O atoms is named using the root of the nonmetal plus suffix –ate Ion w/ fewer O atoms is named using the root of the nonmetal plus suffix –ite Ex: NO3- nitrate NO2- nitrite SO42- sulfate SO32- sulfite Rules for naming halogen-oxyanions: Oxyanion w/ most O atoms is named using prefix per-, root of the nonmetal, and suffix –ate. Oxyanion w/ one less O atom is named using root of the nonmetal and suffix –ate. Oxyanion w/ 2 less O atoms is named using root of the nonmetal and suffix –ite. Oxyanion w/ 3 less O atoms is named using prefix hypo-, root of the nonmetal, and suffix –ite. Ex: ClO4- perchlorate ClO3- chlorate ClO2- chlorite ClO- hypochlorite Some things to remember: Groups 1A and 2A metals have only one oxidation number Transition metals and metals on the right side of the periodic table often have more than one oxidation number Ex: Ions Formula Compound Name Fe2+ & O2- FeO Iron(II) oxide Fe3+ & O2- Fe2O3 Iron(III) oxide METALLIC BONDS AND PROPERTIES OF METALS Metallic bonds – outer energy levels of the metal atoms overlap Electron sea model – all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons Electrons in the outer energy levels of the bonding metallic atoms are not held by any specific atom and can move easily from one atom to the next; these are called delocalized electrons Metallic bond – attraction of a metallic cation for delocalized electrons Properties of metals Melting points vary greatly, but in general metals have high melting and boiling points Ex: Hg is liquid at room temp. 0 Ex: W has a melting point of 3422 C, making it useful for light bulb filaments and spacecraft parts Durable Delocalized electrons in metal are free to move, keeping metallic bonds intact Delocalized electrons move heat from one place to another more quickly than other materials Metal alloys – mixture of elements that has metallic properties Ex: steel is a mixture of iron and at least one other element Substitutional alloy – has atoms of the original metallic solid replaced by other metal atoms of similar size Ex: sterling silver Interstital alloy – form when the small holes in a metallic crystal are filled with smaller atoms Ex: carbon steel – holes in iron crystal are filled with carbon, making the solid harder and stronger THE COVALENT BOND Covalent bond – chemical bond that results from the sharing of valence electrons Generally occurs when elements are relatively close to each other on the periodic table Majority form between nonmetallic elements Molecule – formed when 2 or more atoms bond covalently Single covalent bonds (sigma bonds) Group 7A form a single covalent bond Group 6A share 2 electrons to form 2 covalent bonds Ex: H2O Group 5A form 3 covalent bonds with atoms of nonmetals Ex: F2 Ex: ammonia NH3 Group 4A form 4 covalent bonds Ex: methane CH4 Multiple covalent bonds Atoms of C, N, O, and S most often form multiple bonds O molecule shares 2 electron pairs, forming a double bond: O2 N molecule shares 3 electron pairs, forming a triple bond: N2 Pi bond – parallel orbitals overlap to share electrons Double covalent bond has 1 sigma and 1 pi bond Strength of covalent bonds Bond length – distance between 2 bonding nuclei at the position of maximum attraction As the number of shared electron pairs increases, bond length decreases Shorter bond length makes stronger bond Bond dissociation energy – amount of energy required to break a specific covalent bond As 2 atoms are bonded closer together, more bond energy is needed to separate them NAMING MOLECULES Naming binary molecular compounds – prefixes indicate number of atoms 1-mono 6-hexa 2-di 7-hepta 3-tri 8-octa 4-tetra 9-nona 5-penta 10-deca Exception – first element never uses mono Common names Naming acids – if compound produces H+ ions in solution, it is an acid Binary acids – contains H and one other element Use prefix hydro-, root word of other element, suffix –ic, followed by acid Ex: HBr is hydrobromic acid Oxyacids – contains H and a oxyanion If anion suffix is –ate, it is replaced w/ -ic If anion suffix is –ite, it is replaced w/ -ous Ex: HNO3 is nitric acid HNO2 is nitrous acid Considered acids only when in solution (i.e., dissolved in water) Ex: HCl at STP is hydrogen chloride gas but in solution is hydrochloric acid MOLECULAR STRUCTURES Structural formulas – letter symbols and bonds show relative positions of atoms H is always on the end Central atom is usually closer to the left on the periodic table If central atom is not surrounded by 4 e- pairs, it does not have an octet Must have double or triple bond Ex. problems 9-3 – 9-5 Resonance structures – more than one valid Lewis structure can be written; differ only in position of e- pairs and never atom positions Exceptions to Octet Rule 1. Uneven number of valence electrons - some molecules have an odd number of valence electrons and can’t form an octet 2. Suboctet - some compounds form w/ fewer than 8 electrons present 3. Ex: NO2 Reactive and can share an entire pair of electrons donated by another atom Coordinate covalent bond – when one atom donates a pair of electrons to be shared w/ an atom or ion that needs 2 electrons to become stable Ex: BH3 Expanded octet - some compounds have central atoms that contain more than 8 valence electrons Ex: PCl5 MOLECULAR SHAPE Valence Shell Electron Pair Repulsion (VSEPR) model – used to determine molecular shape Based on arrangement that minimizes repulsion of shared and unshared pairs of electrons around the central atom Bond angle – angle formed by any 2 terminal atoms and the central atom Lone pairs of electrons are not shared between 2 nuclei, therefore they occupy a slightly larger orbital than shared electrons Shared bonding orbitals are pushed together slightly by lone pairs, therefore they occupy a slightly smaller orbital than lone pairs Table 9-3 Hybridization – atomic orbitals are mixed to form new, identical hybrid orbitals Hybrid orbitals – contain one electron that it can share w/ another atom Ex: methane CH4 ELECTRONEGATIVITY AND POLARITY Polar covalent bond – unequal sharing of electrons due to differences in electronegativity Ex: water H2O Microwave ovens Water molecules are constantly moving as they align themselves w/ the changing polarity of the microwave energy field How it Works, p. 270