BONDING General Rule of Thumb: metal + nonmetal = ionic polyatomic ion + metal or polyatomic ion = ionic (both) nonmetal + nonmetal(s) = covalent.
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BONDING
General Rule of Thumb:
metal + nonmetal = ionic
polyatomic ion + metal or polyatomic ion = ionic (both)
nonmetal + nonmetal(s) = covalent
Valence Electrons
Knowing electron configurations is important because the number of valence electrons determines the chemical properties of an element.
Valence Electrons: The e- in the highest occupied energy level of an element’s atoms.
Valence Electrons
All elements in a particular group or family have the same number of valence electrons (and this number is equal to the group number of that element)
Examples: Group 1 elements (Na, K, Li, H): 1 valence electron. Group 2 elements (Mg, Ca, Be): 2 valence electrons. Group 17 elements (Cl, F, Br): 7 valence electrons.
Lewis Structures
Electron dot structures show the valence electrons as dots around the element’s symbol:
Li B Si N O F Ne
Lewis Structures
Electron dot structures show the valence electrons as dots around the element’s symbol:
Li B Si N O F Ne
Octet Rule
Noble gas atoms are very stable; they have stable electron configurations. In forming compounds, atoms make adjustments to achieve the lowest possible (or most stable) energy.
Octet rule: atoms react by changing the number of electrons so as to acquire the stable electron structure of a noble gas.
Octet Rule
Atoms of METALS obey this rule by losing electrons.
Na: Na+: Atoms of NONMETALS obey this rule by
gaining electrons. Cl: Cl-: Transition metals are exceptions to this
rule. Example: silver (Ag) By losing one electron, it acquires a
relatively stable configuration with its 4d sublevel filled (pseudo noble-gas)
Octet Rule
Atoms of METALS obey this rule by losing electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: Cl-: Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled (pseudo noble-gas)
Octet Rule
Atoms of METALS obey this rule by losing electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled (pseudo noble-gas)
Octet Rule
Atoms of METALS obey this rule by losing electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled (pseudo noble-gas)
Octet Rule
Atoms of METALS obey this rule by losing electrons.
Na: 1s22s22p63s1
Na+:1s22s22p6 = same # electrons as Ne
Atoms of NONMETALS obey this rule by gaining electrons.
Cl: 1s22s22p63s23p5
Cl-: 1s22s22p63s23p6
Transition metals are exceptions to this rule. Example: silver (Ag) By losing one electron, it acquires a relatively
stable configuration with its 4d sublevel filled (pseudo noble-gas)
Ionic Bonds
Anions and cations have opposite charges; they attract one another by electrostatic forces (IONIC BONDS)
Ionic Bonds Ionic compounds are electrically
neutral groups of ions joined together by electrostatic forces. (also known as salts) the positive charges of the cations
must equal the negative charges of the anions.
use electron dot structures to predict the ratios in which different cations and anions will combine.
Ex. of Ionic Bonds: Determine charge & CRISS CROSS
Na Cl
Al Br
K O
Mg N
K P
Na+Cl- = NaCl
Al3+Br- = AlBr3
K+O2- = K2O
Mg2+N3- = Mg3N2
K+P3- = K3P
DETERMINE THE CHARGE OF THE ION, CRISS CROSS CHARGES so the compound is NEUTRAL!!! (compounds DO NOT have charges)
Cation + always goes 1st, Anion – always goes last!!
Making Ionic Compounds
+ CATION always goes first in a compound!!!
- ANION always goes last in a compound!!!
Another Example: reducing to the simplest formula
What is the ionic compound formula for the calcium ion with the oxide ion?
Ca O
Ionic Compounds w/ polyatomic ions
POLYATOMIC:DO NOT CHANGE THE FORMULA OF A POLYATOMIC ION!!!
Determine the formula for:
Li+ and PO43-
NH4+ and N3-
Covalent Bonds
Covalent bonds: occur between two or more nonmetals; electrons are shared not transferred (as in ionic bonds)
The result of sharing electrons is that atoms attain a more stable electron configuration.
Covalent Bonds
Most covalent bonds involve: 2 electrons (single covalent bond), 4 electrons (double covalent bond, or 6 electrons (triple covalent bond).
MOLECULES
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
Octet Rule
Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.
THERE ARE A FEW EXCEPTIONS!!!Hydrogen can only have 1 bond (2 electrons around it)
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons)
H2 HBr
CCl4 O2
N2 CO
Writing Lewis Formulas (for molecules) How to:
1. Add up all valence electrons for EACH atom in the molecule
2. Attach atoms with a single bond (skeleton drawing) (C = ALWAYS CENTRAL, H = ALWAYS ON OUTSIDE OF STRUCTURE)
3. Subtract out 2 electrons for each single bond you drew (EACH BOND = 2 electrons)
4. Distribute remaining electrons (in pairs) around atoms to obtain octet rule (except H)
5. If there’s not enough electrons to satisfy the octet rule, make MULTIPLE BONDS (double, triple)
Electronegativity
We’ve learned how valence electrons are shared to form covalent bonds between elements. So far, we have considered the electrons to be shared equally. However, in most cases, electrons are not shared equally because of a property called electronegativity.
Electronegativity
The ELECTRONEGATIVITY of an element is: the tendency for an atom to attract electrons to itself when it is chemically combined with another element.
The result: a “tug-of-war” between the nuclei of the atoms.
Electronegativity
Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value)
*** see Periodic Table, in lower part of element box
Most electronegative element:Fluorine (4.0)
Least electronegative elements: Fr (0.70), Cs (0.79)
Electronegativity Notice the periodic trend:
As we move from left to right across a row, electronegativity increases (metals have low values nonmetals have high values – excluding noble gases)
As we move down a column, electronegativity decreases.
The higher the electronegativity value, the greater the ability to attract electrons to itself.
Nonpolar Bonds
When the atoms in a molecule are the same, the bonding electrons are shared equally.
Result: a nonpolar covalent bond Examples: O2, F2, H2, N2, Cl2
Polar Bonds When 2 different atoms are joined by a
covalent bond, and the bonding electrons are shared unequally, the bond is a polar covalent bond, or POLAR BOND.
The atom with the stronger electron attraction (the more electronegative element) acquires a slightly negative charge.
The less electronegative atom acquires a slightly positive charge.
Predicting Bond Types Electronegativities help us predict
the type of bond:
Electronegativity Difference
Type of Bond Example
0.00 – 0.40 H-H
0.41 – 1.00 H-Cl
1.01 – 2.00 H-F
2.01 or higher Na+Cl-
covalent
(nonpolar)covalent
(slightly polar)covalent
(very polar)
ionic
Polar Molecules A polar bond in a molecule can
make the entire molecule polar
A molecule that has 2 poles (charged regions), like H-Cl, is called a dipolar molecule, or dipole.
Polar Molecules The effect of polar bonds on the polarity of
a molecule depends on the shape of the molecule.
Example: CO2
O = C = O shape: linear
*The bond polarities cancel because they are in opposite directions; CO2 is a nonpolar molecule.
Polar Molecules The effect of polar bonds on the
polarity of a molecule depends on the shape of the molecule.
Water, H2O, also has 2 polar bonds: But, the molecule is bent, so the bonds
do not cancel. H2O is a polar molecule.
Resonance
A molecule or polyatomic ion for which 2 or more dot formulas with the same arrangement of atoms can be drawn is said to exhibit RESONANCE.
Resonance Example CO3
2-
3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3.
Resonance Example CO3
2-
3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3.
Resonance Example CO3
2-
3 resonance structures can be drawn for CO32-
the relationship among them is indicated by the double arrow.
the true structure is an average of the 3
Resonance Structures
Another way to represent this is by delocalization of bonding electrons:
(the dashed lines indicate the 4 pairs of bonding electrons are equally distributed among 3 C-O bonds; unshared electron pairs are not shown)
Molecular Shape Lewis structures (electron dot
structures) show the structure of molecules…but only in 2 dimensions (flat).
BUT, molecules are 3 dimensional! for example, CH4 is:
Molecular Shape Lewis structures (electron dot structures)
show the structure of molecules…but only in 2 dimensions (flat).
BUT, molecules are 3 dimensional! but in 3D it is:
a tetrahedron!= coming out of
page= going into page= flat on page
Why do molecules take on 3D shapes instead of being flat?
Valence Shell Electron Pair Repulsion theory
“because electron pairs repel one another, molecules adjust their shapes so that the valence electron pairs are as far apart from another as possible.”
Why do molecules take on 3D shapes instead of being flat?
Valence Shell Electron Pair Repulsion theory
Remember: both shared and unshared electron pairs will repel one another (unshared electron pairs repel MORE than shared electron pairs in bonds)
H—N — H—
H
Non-BondingPairs
BondingPairs
5 Basic Molecule Shapes
2. Pyramidal
Example: NH3
(note: unshared pair of electron repels, but is not considered part of overall shape; no atom there to contribute to the shape)
CHEAT SHEET…Let’s make this easy to remember shapesLinear= only 2 regions of SHARED electron pairs
(no unshared electrons) around the CENTRAL atom
Bent = 2 SHARED electron pairs, 2 UNSHARED electron pairs around the CENTRAL atom
Tetrahedral = 4 SHARED electron pairs (meaning 4 bonds) around the CENTRAL atom
Trigonal Planar = 3 SHARED electron pairs (meaning 3 bonds) around the CENTRAL atom
Pyramidal = 3 SHARED electron pairs, 1 UNSHARED pair of electrons around the CENTRAL atom
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