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BONDING AND GEOMETRY. Unit 10 Chemistry Langley. **Corresponds to Chapter 7 and 8 (pages 186-247) in the Prentice Hall Chemistry textbook. PERIODIC TABLE REVIEW. Location of Metals and Nonmetals on the periodic table: Metals are to the left of the “staircase”
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BONDING AND GEOMETRY Unit 10 Chemistry Langley **Corresponds to Chapter 7 and 8 (pages 186-247) in the Prentice Hall Chemistry textbook
PERIODIC TABLE REVIEW • Location of Metals and Nonmetals on the periodic table: • Metals are to the left of the “staircase” • Nonmetals are to the right of the “staircase” • For bonding, the 7 metalloids will treated as metals • All though hydrogen is to the left of the “staircase”, it is not, nor has it ever been a metal. IT IS A NONMETAL!
ATOMS AND IONS REVIEW • Atoms are neutral • They have the same number of protons and electrons • Number of positives = number of negatives • Example: Na 11 protons, 11 electrons 11 – 11 = 0 • Ions have a charge • They have a different number of protons and electrons • Example: Na+111 protons, 10 electrons 11 – 10 = +1 • If an atom GAINS an electron becomes negatively charged ANION • If an atom LOSES an electron becomes positively charged CATION
TYPES OF BONDS • Bonding occurs because every element is either trying to get to 0 electrons in the valence or 8 electrons in the valence (zero and 8 are both stable) • Valence is the outer electron shell—place where bonding occurs • Ionic – Bonding between a metal and a nonmetal • Metallic – Bonding between two metals • Covalent – Bonding between two nonmetals
IONIC BONDING • Very stable and strong • Strongest possible bond • Requires a large amount of energy to break an ionic bond • Forms compounds known as “ionic compounds” • All ionic compounds will dissolve in water and carry a current (electrolyte) • Generally have high melting and boiling points • Compounds are generally hard and brittle
IONIC BONDING • Draw the dot diagram for Na AND Cl Na has 1 valence electron, wants to give that 1 away and get to zero and be stable Cl has 7 valence electrons, wants to get 1 electron so it can get to eight and be stable Na give an electron to Cl and Cl takes that electron from Na
METALLIC BONDING • Metal atoms are pieces of metal that consist of closely packed cations (positively ions) • Cations are surrounded by mobile valence electrons that are free to drift from one part of the metal to another • Metal atoms are arranged in very compact and orderly (crystalline) patterns • Metallic bonding is the electrostatic attraction between conduction electrons, and the metallic ions within the metals, because it involves the sharing of free electrons among a lattice of positively-charged metal ions • Occurs between 2 or more metals • Result of the attraction of free floating valence electrons for the positive ion • These bonds hold metals together
METALLIC BONDING • Properties of metallic bonds • Good conductors of electricity • Electrons are free flowing • Malleablehammered into sheets • Ductiledrawn into wires • Alloy-two metals are bonded together to get the benefits of each • 14 karat gold
COVALENT BONDING • Covalent: • Prefix “co” means share, together • “valent” means valence • Covalent bonds are when atoms SHARE VALENCE electrons • A covalent compound is called a molecule • Covalent bond ALWAYS occurs between 2 nonmetals
TYPES OF COVALENT BONDS • Single Bond • Covalent bond where one pair of electrons (2 electrons total) are shared between 2 atoms • Atoms share electrons so that each has a full octet (8 valence) • Electrons that are shared count as valence electrons for both atoms • Examples • HCl • Cl2
COVALENT BONDING • Double Bonds • Bond in which two pairs of electrons (4 electrons total) are shared between 2 atoms • Examples • O2 • C2F2 • Triple Bonds • Bond in which 3 pairs of electrons (6 total electrons) are shared between atoms • Examples • N2 • AsP
COVALENT BONDING • Covalent Bonds with more than 2 atoms • Examples • CH4 • OF2 • Electron Pairs • Electron pairs involved in the actual bond are called BONDING PAIR or SHARED PAIR electrons • Electrons not involved in the actual bond, those surrounding the rest of each element are called LONE PAIR electrons
POLAR BONDS AND MOLECULES • Covalent bonds are formed by sharing electrons between two atoms • The bonding pair of electrons is shared between both elements, but each atom is tugging on the bonding pair • When atoms in a molecule are the same (diatomic) the bonding pair is shared equallythis bond is called non polar covalent • When atoms in a molecule are different, the bonding pair of electrons are not shared equallythis is called a polar covalent bond
POLAR BONDS AND MOLECULES • Why is the bonding pair not shared equally? • The answer lies within electronegativity • One of the elements is more electronegative than the other and therefore has a greater desire for the shared pair • The MORE electronegative element tends to pull the electrons closer and thus has a slightly negative charge • The LESS electronegative element has a slightly positive charge since the shared pair is being pulled away
POLAR BONDS AND MOLECULES • Drawing/Indicating Polarity
POLAR BONDS AND MOLECULES • Polar Molecules • Molecule in which one end of the molecule is slightly negative and the other end is slightly positive • Just because a molecule contains a polar bond DOES NOT mean the entire molecule is polar • The effect of polar bonds on the polarity of an entire molecule depends on the shape of the molecule and the orientation of the polar bonds
POLAR BONDS AND MOLECULES • Example: CO2O = C = O • Carbon and Oxygen lie along the same axis. • Bond polarities are going to cancel out because they are in opposite directions • Carbon dioxide is a nonpolar molecule even though there are two polar bonds present • Would cancel out if the polarities moved towards each other as well • When polarities cancel out, the molecule is non-polar
POLAR BONDS AND MOLECULES • Example: H2O • Example: CH3Br
FORCES IN A MOLECULE • Dipole-Dipole Forces • Dipoles are created when equal but opposite charges are separated by a short distance • Have to have a positive and a negative end so that one of the elements is pulling on the electron • Only happens in polar molecules • Dipole forces are extremely strong and lead to high melting and boiling points
FORCES IN A MOLECULE • Hydrogen Bonding • Very strong type of dipole force • Only occurs when hydrogen is covalenty bonded to a highly electronegative atom • Always involves hydrogen • Example: HF, HCl
FORCES IN A MOLECULE • London Dispersion Forces • Electrons are in constant motion around a nucleus • At any given time there might be more electrons on one side of an atom than on the other • For a split second, the side with more electrons is negative, and the side with less electrons is positive
FORCES IN A MOLECULE • London Dispersion Forces • Recall that Noble Gases have a full outer shell and you have been told they are unreactive BUT due to London Dispersion Forces, they COULD bond for an instant • Example: Ar2 • London Forces are very weak • The smaller the mass of the atom, the smaller the London Force
BOND DETAILS • Terminology • Bond strength-energy required to break a bond • Bond axis-imaginary line joining two bonded atoms (example: C-C) • Bond length-the distance between two bonded atoms at their minimum potential enery; the average distance between two bonded atoms • Bond energy-energy required to break a chemical bond and form neutral isolated atoms • Chemical compound tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level
BOND DETAILS • Comparison of Bond Length/Strength for Covalent Bond Types: • Longer bond = less bond strength • Rating 1-3 (with 3 as the largest and 1 as the smallest)
BOND DETAILS • Coordinate Covalent Bonds • Very rare • Tend to form harmful molecules • Occurs when both of the bonding pair of electrons in a covalent bond come from only ONE of the atoms • Example: CO
BOND DETAILS • Resonance • Occurs when there are more than one possible structures for a molecule • Refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • Example: CO2 • To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures • Even though all of the structures are different, the number of bonding pair of electrons and lone pair of electrons stay the same in each structure
VSEPR THEORY • Valence Shell Electron Pair Repulsion Theory • Allows us to picture molecules in 3 dimensions • Centers around the fact that electrons have negative charges and repel one another • So electron pairs within a structure try to arrange themselves to be as far away from other pairs as possible
VSEPR THEORY • Tetrahedral • Central atom bonds to 4 atoms and has zero lone pairs • CH4
VSEPR THEORY • Pyramidal • The central atom bonds to 3 atoms and has 1 lone pair of electons • NH3
VSEPR THEORY • Trigonal Planar • The central atom bonds to 3 atoms and has zero lone pairs • CO3-2
VSEPR THEORY • Bent Triatomic • The central atom bonds to 2 atoms and has 2 lone pair of electrons • H2O
VSEPR THEORY • Linear Triatomic • The central atom bonds to 2 atoms and has zero lone pair of electrons • CO2
VSEPR THEORY • Linear • One bond between 2 atoms • HCl • N2