BONDING AND GEOMETRY Unit 10 ChemistryLangley **Corresponds to Chapter 7 and 8 (pages ) 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 +1  11 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  Malleable  hammered into sheets  Ductile  drawn 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  Cl 2

COVALENT BONDING  Double Bonds  Bond in which two pairs of electrons (4 electrons total) are shared between 2 atoms  Examples  O 2  C 2 F 2  Triple Bonds  Bond in which 3 pairs of electrons (6 total electrons) are shared between atoms  Examples  N 2  AsP

COVALENT BONDING  Covalent Bonds with more than 2 atoms  Examples  CH 4  OF 2  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 equally  this bond is called non polar covalent  When atoms in a molecule are different, the bonding pair of electrons are not shared equally  this 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: CO 2 O = 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: H 2 O  Example: CH 3 Br

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: Ar 2  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) BondLengthStrength Single31 Double22 Triple13

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: CO 2  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  CH 4

VSEPR THEORY  Pyramidal  The central atom bonds to 3 atoms and has 1 lone pair of electons  NH 3

VSEPR THEORY  Trigonal Planar  The central atom bonds to 3 atoms and has zero lone pairs  CO 3 -2

VSEPR THEORY  Bent Triatomic  The central atom bonds to 2 atoms and has 2 lone pair of electrons  H 2 O

VSEPR THEORY  Linear Triatomic  The central atom bonds to 2 atoms and has zero lone pair of electrons  CO 2

VSEPR THEORY  Linear  One bond between 2 atoms  HCl  N 2