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Periodic Trends Bellringer
1) Which has a larger size, a Mg atom or a Mg ion? 2) smaller ionization energy, K or Br? 3) smaller size, F atom or an I atom? 4) smaller electronegativity, O or Se? 5) less shielding, Ca or Sr?
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Answers 1) Mg atom 2) K 3) F 4) Se 5) Ca
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Ch. 7 Ionic and Metallic Bonding
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Valence Electrons Valence electrons = electrons in the highest occupied energy level of an element’s atoms This # largely determines the chem. properties of the element To find the # of valence e-’s in an atom of a representative element, look at its group # Group 1 has 1, group 2 has 2, group 13 or 3A has 3, … Electron dot structures (Lewis dot diagrams) = diagrams that show valence electrons as dots *Draw Lewis dot structures*
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Octet Rule Octet rule = atoms gain or lose electrons to achieve a stable level of usually 8 atoms of metals tend to lose their valence electrons leaving a complete octet in the next-lowest energy level Atoms of some nonmetals tend to gain electrons or to share electrons w/ another nonmetal to achieve a complete octet
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Cations Cations (+ charged atom) are formed when they lose one or more valence electrons (ionization) in order to become stable Usually lose from 1 to 3 valence electrons tops Can use electron configurations to illustrate the point but let’s use electron dot diagrams for ease Exceptions are due to having to lose or gain too many valence electrons to achieve a noble gas state, so: Some atoms attain a pseudo noble-gas electron configurations – Cu - copper, Ag - silver, Au - gold, Cd - cadmium Two atoms are walking down the street. Says one atom to the other, "Hey! I think I lost an electron!" The other says, "Are you sure??" "Yes, I'm positive!"
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Anions Anions (- charged atoms) are formed when they gain 1 or more valence electrons Typically suffix is –ide Halide ions (F-, Cl-, Br-, I-) are halogens that gain 1 e- CP 7.1, PP 1-2 pg. 193, sect. assessment 7.1 pg
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Common ions Cation Name Anion Name* H 1+ hydrogen H 1- hydride
Li 1+ lithium F 1- fluoride Na 1+ sodium Cl 1- chloride K 1+ potassium Br 1- bromide Cs 1+ cesium I 1- iodide Be 2+ beryllium O 2- oxide Mg 2+ magnesium S 2- sulfide Al 3+ aluminum Ag 1+ silver *The root is given in color.
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Ionic Compounds Anions and cations are held together by opposite charges. Ionic compounds are called salts Electrically neutral Simplest ratio is called the formula unit The bond is formed through the transfer of electrons (called an ionic bond) Electrons are transferred to achieve noble gas configuration Most are crystalline solids at room temp. High melting points – large attractive forces result in a very stable structure Good conductors of electricity when melted or dissolved in water Overheard at the mall Teen 1: Did you hear oxygen and magnesium got together?? Teen 2: OMg! Your mama's so ugly Your mama's so ugly...even Fluorine won't bind to her
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Formulas Chemical formula = shows the kinds and #s of atoms in the smallest representative unit of a substance NaCl is the chem. formula for sodium chloride Formula unit = the lowest whole-number ratio of ions in an ionic compound MgCl2, NaCl, AlBr3 CP 7.2, PP pg. 196, 7.2 sect. assessment pg
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Ch. 7 Bellringer Write the Lewis electron-dot symbol for each of the following A) sodium B) fluorine C) magnesium ion (Mg2+) Write the chemical formula that results when the following pairs of ions combine to form an ionic bond D) Mn4+ and O2- E) Li1+ and Cl1-
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Metallic Bonds and Properties
Metallic bonds = consist of the attraction of the free-floating valence electrons for the positively charged metal ions How atoms are held together in the solid Metals hold onto their valence electrons very weakly Think of them as positive ions floating in a “sea of electrons” Electrons are free to move through the solid Metals conduct electricity Malleable - hammered into shape (bend) Ductile - drawn into wires Electrons allow cations to slide by each other under pressure Metals are crystalline structures and atoms are arranged in very compact and orderly patterns *Cu vs. Cu compounds hammer demo – pg. 202*
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Alloys Alloys = mixtures composed of 2 or more elements, at least one of which is metal Brass – copper and zinc Sterling silver – silver and copper Bronze – copper and tin Steel – Fe, Cr, and others Impt. b/c their properties are often superior to those of their component elements (usually cheaper as well) *Making an alloy DEMO* - pg. 205
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Ch. 8 Covalent Bonding
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Molecules Covalent bond = formed by sharing electrons b/w 2 or more atoms Molecule = a neutral group of atoms joined by covalent bonds Diatomic molecule = a molecule consisting of 2 of the same atoms H2, N2, O2, F2, Cl2, Br2, I2 are the diatomic molecules in nature Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds Most molecular compounds are composed of atoms to 2 or more nonmetals *Make a table comparing covalent bonding and ionic bonding*
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Molecular Formula Molecular formula = the chemical formula for a molecular compound Shows how many atoms of each element a molecule contains H2O, CO2, C2H6, O2 8.1 Sect. assessment, Pg
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Covalent Bonding Electron sharing usually occurs so that atoms attain the electron configuration of noble gases Combos of atoms of the nonmetals and metalloids in 4A, 5A, 6A, and 7A are likely to form covalent bonds Single covalent bond = 2 atoms sharing 1 pair of electrons H2 2 dots in an electron dot diagram represents this bond A dash in a structural formula represents this bond A molecular formula does NOT show this bond only the # of atoms Halogens form these bonds in their diatomic molecules Unshared pair = pair of valence electrons not shared in an electron dot diagram *CP 8.1, PP 7-8 pg. 220
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Double and Triple Covalent Bonds
Double covalent bond = involves 2 shared pairs of electrons oxygen Triple covalent bond = involves 3 shared pairs of electrons nitrogen
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Coordinate covalent bonds
The shared electron pair comes from one of the bonding atoms CO – look at pg. 223 Polyatomic ion = a tightly bound group of atoms that has a positive or negative charge and behaves as a unit NH4+ *c.p. 8.2, p.p pg. 225 A sign outside the chemistry hotel reads "Great Day Rates, Even Better NO3-'s"
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Bond Dissociation Energies
Bond dissociation energy = the energy required to break the bond b/w 2 covalently bonded atoms a large bond dissociation energy corresponds to a strong covalent bond H2 = 435 kJ/mol, C-C single bond = 347 kJ/mol, C=C double bonds = 657 kJ/mol, and triple bonds = 908 kJ/mol
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Resonance Structure Resonance structure = a structure that occurs when it is possible to draw 2 or more valid electron dot structures that have the same # of electron pairs for a molecule or ion Double-headed arrows are used to connect Double bonds are usually shorter than single bonds but they are the same lengths b/c it is an avg. of the 2 structures resonance
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Exceptions to the Octet Rule
Cannot be satisfied in molecules whose total # of valence electrons is an odd #. NO2, ClO2, NO Sometimes w/ an even # as well (Fewer or more) BF3 PCl5 SF6
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*8.2 sect. assessment pg. 229* *Exceptions to the octet rule: A resonance hybrid teacher DEMO*
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Molecular Orbitals Molecular orbitals = orbitals that apply to the entire molecule just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole Bonding orbital = a molecular orbital that can be occupied by 2 electrons of a covalent bond
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Sigma Bonds Sigma bonds = formed when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting 2 atomic nuclei Atomic orbitals overlap end to end Two s orbitals can combine to form a molecular orbital H2 Two p orbitals F2 The attractions b/w electrons and nuclei of two atoms overpower the repulsions b/w the 2 nuclei or b/w the 2 sets of electrons = covalent bond (stable molecule)
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Pi Bonds Pi bond = a covalent bond in which the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms Orbitals overlap side by side Atomic orbitals in pi bonding overlap less than in sigma bonding – weaker than sigma bonds A typical double bond consists of 1 sigma & 1 pi bond, triple bond is 1 sigma & 2 pi bonds In special cases, they form w/o any sigma bonds
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VSEPR Theory VSEPR theory (valence-shell electron-pair repulsion theory) = the repulsion b/w electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible Explains the actual 3-D shapes of molecules
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Hybrid Orbitals Hybridization = several atomic orbitals mix to form the same total number of equivalent hybrid orbitals Single bonds – one 2s and three 2p orbitals mix to form four sp3 hybrid orbitals Double bonds – one 2s and two 2p orbitals mix to form three sp2 hybrid orbitals Triple bonds – one 2s and 1 2p orbitals mix to form two sp hybrid orbitals
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Electronic Geometry 2 electron densities – Linear
3 e- densities – Trigonal Planar 4 e- densities – Tetrahedral 5 e- densities – Trigonal Bipyramidal 6 e- densities – octahedral
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Linear Atoms connected in a straight line
All molecules w/ 2 atoms and some w/ 3 180⁰ bond angle HCl, CO2 Hybridization – sp 2 bonds/0 lone pairs
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Trigonal Planar Triangular flat 120⁰ bond angle BCl3
Hybridization – sp2 3 bonds/0 lone pairs Bent – 2 bonds/1 lone pair 118⁰ SO2
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Tetrahedral 4 surfaces 109.5⁰ bond angle CH4 Hybridization – sp3
4 bonds/0 lone pairs Pyramidal (trigonal pyramidal) – 3 bonds/1 lone pair 107⁰ NH3 Bent – 2 bonds/2 lone pairs 105⁰ H2O
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Trigonal Bipyramidal 90, 120, and 180⁰ bond angles PF5
Hybridization – sp4 5 bonds/0 lone pairs See-saw – 4 bonds/0 lone pairs 90, 120, and 180⁰ SF4 Tee-shaped – 3 bonds/2 pairs 90 and 180⁰ ClF3 Linear – 2 bonds/3 lone pairs 180⁰ XeF2
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Octahedral 90 and 180⁰ bond angles SF6 Hybridization – sp5
6 bonds/0 lone pairs Square pyramidal – 5 bonds/1 lone pair 90 and 180⁰ BrF5 Square Planar – 4 bonds/2 lone pairs XeF4 8.3 sect. assessment pg
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The VSEPR Model .. .. .. The Shapes of Some Simple ABn Molecules O S O
Linear Bent Trigonal planar Trigonal pyramidal AB6 F P F S F Cl F Xe Students often confuse electron-domain geometry with molecular geometry. You must stress that the molecular geometry is a consequence of the electron domain geometry. The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them. T-shaped Square planar Trigonal bipyramidal Octahedral
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Bond Polarity Nonpolar covalent bonds = equal sharing of electrons
H2, O2, N2, Cl2 Polar covalent bond = unequal sharing of electrons HCl, H2O The more electronegative atom attracts electrons more strongly and gains a slightly negative charge, the less electronegative atom has a slightly positive charge
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Table 8.3 pg. 238 C.P. 8.3, P.P pg. 239
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Polar Molecules Polar molecule = one end of the molecule is slightly negative, the other is slightly positive Dipole = a molecule that has 2 poles w/opposite charges The shape of a molecule and the polarity of its bonds together determine whether the molecule is polar or nonpolar Equal and opposite directions arrows cancel = nonpolar Arrows same direction = polar
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Attractions b/w Molecules
Weaker than ionic or covalent bonds Van der Waals forces = weak intermolecular attractions 1. Dipole interactions = intermolecular forces resulting from the attraction of oppositely charged regions of polar molecules 2. Dispersion forces = (weakest of all) attractions b/w molecules caused by the electron motion of one molecule affecting the electron motion on the other through electrical forces Hydrogen bonds = attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom The combo of this strongly polar bond and the lack of shielding effect in a hydrogen atom is responsible for the relative strength of hydrogen bonds (strongest of all intermolecluar attractions)
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Table 8.4 pg. 244 *8.4 section assessment pg. 244
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