Bonding
This presentation is the property of its rightful owner.
Sponsored Links
1 / 68

BONDING PowerPoint PPT Presentation


  • 79 Views
  • Uploaded on
  • Presentation posted in: General

BONDING. General Rule of Thumb: metal + nonmetal = ionic polyatomic ion + metal or polyatomic ion = ionic (both) nonmetal + nonmetal(s) = covalent. Ch 8 & 9 – Honors Chemistry. Ionic Bonds. Isn’t it ionic that opposites attract?. Valence Electrons.

Download Presentation

BONDING

An Image/Link below is provided (as is) to download presentation

Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author.While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server.


- - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - -

Presentation Transcript


Bonding

BONDING

General Rule of Thumb:

metal + nonmetal = ionic

polyatomic ion + metal or polyatomic ion = ionic (both)

nonmetal + nonmetal(s) = covalent

Ch8 & 9 – Honors Chemistry


Ionic bonds

Ionic Bonds

Isn’t it ionic that opposites attract?


Valence electrons

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 electrons1

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

Lewis Structures

  • Electron dot structures show the valence electrons as dots around the element’s symbol:

  • Li

  • B

  • Si

  • N

  • O

  • F

  • Ne


Lewis structures1

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

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 rule1

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 rule2

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 rule3

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 rule4

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)


Ionic bonds1

Ionic Bonds

  • Anions and cations have opposite charges; they attract one another by electrostatic forces (IONIC BONDS)


Ionic bonds2

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.


Examples of ionic bonds

Examples of Ionic Bonds

Na+Cl- = NaCl

  • NaCl

  • AlBr

  • K O

  • MgN

  • KP

Al3+Br- = AlBr3

K+O2- = K2O

Mg2+N3- = Mg3N2

K+P3- = K3P


Criss cross method

Criss Cross Method

  • Criss-cross ionic charges down as subscripts (without +/-) and reduce (to determine lowest whole # ratio of cation:anion)

  • ANY TIME YOU ADD A SUBSCRIPT TO A POLYATOMIC ION, YOU MUST FIRST PUT THAT ION IN PARENTHESES.


Covalent bonds

Covalent Bonds

The joy of sharing!


Covalent bonds1

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 bonds2

Covalent Bonds

  • Most covalent bonds involve:

    • 2 electrons (single covalent bond),

    • 4 electrons (double covalent bond, or

    • 6 electrons (triple covalent bond).


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Bonding

  • 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)

  • H2HBr

  • CCl4O2

  • N2CO


Octet rule5

Octet Rule

  • Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.


Writing lewis structures

Writing Lewis Structures

  • Select a skeleton for the molecule (the least electronegative element is usually the central element).

  • Calculate N (the # of valence e- need by all atoms in the molecule of polyatomic ion.

  • Calculate A (the # of electrons available).

  • Calculate S (the # of electrons shared in the molecule) S = N – A

  • Place the S electrons as shared pairs in the skeleton.

  • Place the additional electrons as unshared pairs to fill the octet of every representative elements (except hydrogen!).


Lewis structure examples

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples1

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples2

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples3

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples4

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples5

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples6

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Lewis structure examples7

CO2-

OH-

NO3-

SO42-

CBr3-

N22-

CO32-

NH4+

Lewis Structure Examples:


Electronegativity

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.


Electronegativity1

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.


Electronegativity2

Electronegativity

  • Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value)

  • **See Figure 6-18 p. 169 (Honors)

    • Most electronegative element:Fluorine (3.98)

    • Least electronegative elements:

      Fr (0.70), Cs (0.79)


Electronegativity3

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

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

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.


Polar bonds1

Polar Bonds

  • Example: HCl

  • Electronegativities:

    • H = 2.20

    • Cl = 3.16

-

+

H

Cl


Polar bonds2

Polar Bonds

  • Example: H2O

  • Electronegativities:

    • H = 2.20

    • O = 3.44


Polar bonds3

Polar Bonds

  • Example: H2O

  • Electronegativities:

    • H = 2.20

    • O = 3.44


Polar bonds4

Polar Bonds

  • Example: H2O

  • Electronegativities:

    • H = 2.20

    • O = 3.44


Predicting bond types

Predicting Bond Types

  • Electronegativities help us predict the type of bond:

covalent

(nonpolar)

covalent

(slightly polar)

covalent

(very polar)

ionic


Polar molecules

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 molecules1

Polar Molecules

  • The effect of polar bonds on the polarity of a molecule depends on the shape of the molecule.

  • Example:CO2

    O = C = Oshape: linear

    *The bond polarities cancel because they are in opposite directions; CO2 is a nonpolar molecule.


Polar molecules2

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

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

Resonance Example

  • CO32-

  • 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 example1

Resonance Example

  • CO32-

  • 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 example2

Resonance Example

  • CO32-

  • 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

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)

    • See p. 256


Vsepr

VSEPR

valence shell electron pair repulsion


Molecular shape

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 shape1

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

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 flat1

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.

Non-Bonding

Pairs

H—N — H

Bonding

Pairs

H


5 basic molecule shapes

5 Basic Molecule Shapes

  • Linear

  • Example: CO2


5 basic molecule shapes1

5 Basic Molecule Shapes

  • Bent or angular

  • Example: H2O

  • Notice electron pair repulsion


5 basic molecule shapes2

5 Basic Molecule Shapes

  • tetrahedral

  • example: CH4


5 basic molecule shapes3

5 Basic Molecule Shapes

  • 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)


5 basic molecule shapes4

5 Basic Molecule Shapes

  • Trigonal planar or planar triangular

  • Example: BF3


Geometry and polarity

Geometry and polarity

  • Three shapes will cancel out polarity.

  • Shape One: Linear


Geometry and polarity1

Geometry and polarity

  • Three shapes will cancel out polarity.

  • Planar triangles

120º


Geometry and polarity2

Geometry and polarity

  • Three shapes will cancel out polarity.

  • Tetrahedral


Geometry and polarity3

Geometry and polarity

  • Others don’t cancel

  • Bent


Geometry and polarity4

Geometry and polarity

  • Others don’t cancel

  • Trigonal Pyramidal


  • Login