Bettelheim, Brown, Campbell and Farrell Chapter 3 Chemical Bonds Bettelheim, Brown, Campbell and Farrell Chapter 3

Ionization Energy Energy required to remove outermost electron (most loosely held)

Noble Gas Configurations Noble gas configuration s2p6 very stable

The Octet Rule Octet rule: Group 1A-7A elements to achieve an outer shell of eight valence electrons Anion: Negative ion formed when an atom gains electrons Cation: Positive ion formed when an atom loses electrons

The Octet Rule—Cations Cation: Sodium atom loses an electron to form a sodium ion, which has the same electron configuration as neon Na (11 electrons): 1s2 2s2 2p6 3s1 Na+ (10 electrons): 1s2 2s2 2p6

The Octet Rule—Anions Anion: Chlorine atom gains an electron to form a chloride ion, which has the same electron configuration as argon chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5 chloride ion (18 electrons): 1s2 2s2 2p6 3s2 3p6

Cation Names Groups 1A, 2A, and 3A The name of the element followed by the word “ion”

Transition Metal Cations Cations derived from transition and inner transition elements more than one type of cation Stock System (IUPAC): Use Roman numerals to show charge: Fe2+ is Iron (II) Fe3+ is Iron (III) Cu+ is Copper (I) Cu2+ is Copper (II) Old System: Use the suffix -ous to show the lower positive charge and the suffix -ic to show the higher positive charge Fe2+ is Ferrous Fe3+ is Ferric Cu+ is Cuprous Cu2+ is Cupric

Transition Metal Ion Names

Anion Names Add “ide” to the root name of the element

Polyatomic Ions Contain two or more atoms Common names often used (in parentheses)

Naming Ionic Compounds Name the positive ion first, then the negative ion Number of each ion not included NaBr Al2O3 MgSO4 K2S (NH4)3PO4

Naming Ionic Compounds NaBr Sodium bromide Al2O3 Aluminum oxide MgSO4 Magnesium sulfate K2S Potassium sulfide (NH4)3PO4 Ammonium phosphate

Formulas of Ionic Compounds The total number of positive charges must equal the total number of negative charges Li+ and Br- form LiBr (+1) + (-1) Ba2+ and I- form BaI2 (+2) + 2(-1) Al3+ and S2- form Al2S3 2(+3) + 3(-2)

Forming Chemical Bonds Ionic bond: the force of electrostatic attraction between a cation and an anion Atom loses or gains electrons to make a filled valence shell (octet) and become an ion. Covalent bond: a pair of electrons that are shared by two atoms Atom shares electrons to make a filled valence shell (octet)

Forming an Ionic Bond--NaCl Formation of sodium chloride, NaCl Single-headed curved arrow used to show the transfer of the electron

Ionic Bonds Force of attraction between a cation and an anion. Depends on electronegativity measure of an atom’s attraction for shared pair of electrons in chemical bond with another atom)

Covalent Bonds Result of one or more pairs of electrons that are shared by two atoms Each atom has full valence shell (octet) In H2, each hydrogen contributes one electron to the single bond

Molecular Compounds Molecular compound: only covalent bonds Naming molecular compounds the less electronegative element is named first (it is generally written first in the formula) prefixes “di-”, tri-”, etc. are used to show the number of atoms of each element; the prefix “mono-” is generally omitted Exception: carbon monoxide NO is nitrogen oxide (nitric oxide) SF2 is sulfur difluoride N2O is dinitrogen oxide (laughing gas)

Electronegativity F has highest value Noble gases have 0 value Ionic bonds form when electronegativity difference ≥ 1.9

Polarity of Bonds Nonpolar: Electrons are shared equally Polar: Electrons are NOT shared equally

Polarity of Covalent Bonds

Polarity of Covalent Bonds More electron density shown by δ- or the head of a crossed arrow Less electron density shown by δ+ or the tail of a crossed arrow

Polarity of Molecules Polar molecule has polar bonds, and Has partial positive and partial negative charges in different parts of molecule, i.e., is a dipole (has two poles) Carbon dioxide, CO2, has two polar bonds but, because of its geometry, the pulls balance out so it is a nonpolar molecule

Polarity of Molecules Water, H2O, has two polar bonds and, because of its geometry, is a polar molecule

Polarity of Molecules Both dichloromethane, CH2Cl2, and formaldehyde, CH2O, have polar bonds and are polar molecules

Lewis Structures Used to decide on the arrangement of atoms in the molecule Bonding (shared) electrons are shown as bonds (lines) Nonbonding electrons are represented as a pair of Lewis dots

Drawing Lewis Structures 1. Determine the number of valence electrons in the molecule 2. Decide on the arrangement of atoms in the molecule 3. Connect the atoms by single bonds 4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots 5. In a single bond, atoms share one pair of electrons; in a double bond, they share two pairs, and in a triple bond they share three pairs.

Exceptions to the Octet Rule H and He have a maximum of 2 electrons (duet) Period 2 elements have a maximum of 8 electrons (use 2s and 2p orbitals) Atoms of period 3 elements may have more than 8 electrons

Lewis Structures Examples: (the number of valence electrons is given in parentheses after the molecular formula

Lewis Structures Examples NH3 CH3OH CH3COOH

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Valence-Shell Electron-Pair (VSEPR) Model Because like charges repel each other, the various regions of electron density around an atom spread so that they are as far away from each other as possible CH4: measured H-C-H bond angles are 109.5°

Summary of Molecular Geometry # Electron Lone Bond Angle Shape regions pairs 2 0 180o Linear 3 0 120o Planar 2 1 Angular (bent) 4 0 109.5o Tetrahedral 3 1 Trigonal pyramidal 2 2 Angular (bent)

VSEPR Model the measured H-N-H bond angles are 107.3° the unshared pair is not shown on this model the measured H-O-H bond angle is 104.5°