Unit 2 Bonding Valence electrons are the outer
Unit 2 - Bonding
Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. Group e- configuration # of valence e- I ns 1 1 II ns 2 2 III ns 2 np 1 3 IV ns 2 np 2 4 V ns 2 np 3 5 VI ns 2 np 4 6 VII ns 2 np 5 7 9. 1
Lewis Dot Symbols for the Representative Elements & Noble Gases 9. 1
Trends in the Periodic Table �The elements have a regular (periodic) recurrence of physical and chemical properties �The horizontal rows are called periods �The vertical columns are called groups
Atomic Size �Covalent Atomic Radius – half of the distance between the nuclei of two atoms in a homonuclear diatomic molecule �Group trend – atomic size increases as we move down a group �The farther away the e- from the nucleus, the less strongly they are held �Periodic trend – atomic size decreases as we move from left to right �Nuclear charge increases as we move left to right, pulling the e- closer to the nucleus
Ionic Size �Cations are always smaller than neutral atoms because they have less e�Loss of outer shell e- results in increased attraction by the nucleus for the remaining e�Anions are always larger than neutral atoms because they have more e�Addition of outer shell e- results in less attraction to the protons of the nucleus
Ionization Energy �The energy required to remove an electron from a gaseous atom Na (g) + Energy Na+(g) + e�Group trend – Ionization energy decreases as we move down a group because the size of the atom is increasing, with more orbitals of e-, allowing the outermost e- to be easily removed �Periodic trend – Ionization energy increases from left to right. Nuclear charge is increasing so more energy is required to remove e-
Electron Affinity �The energy change (release) that accompanies the addition of an electron to a gaseous atom F (g) + e- F-(g) + Energy �Group trend – Electron affinity generally decreases with increasing atomic size (releases less energy as you go down) �Periodic trend – Generally increases from left to right because atoms become smaller due to increased nuclear charge (releases more energy as you go left to right)
Electronegativity �The tendency for an atom to attract electrons to itself. It is affected by the distance from the atom’s valence electrons to the nucleus. �Group trend – Electronegativity usually decreases as you move down a group �Periodic trend – Electronegativity usually increases as you move across a period �Do the Noble Gases have electronegativity?
�Why do substances bond? �More stability �Atoms want to achieve a lower energy state
Chemical bonds: an attempt to fill electron shells 1. 2. 3. Ionic bonds – Covalent bonds – Metallic bonds
Ionic Bonding �Between a metal and a non-metal with very different electronegativity. �Metals lose electrons becoming a cations, while nonmetals gain electrons becoming anions. �An ionic bond is an electrostatic attraction between the oppositely charged ions.
Ionic Bonds: One Big Greedy Thief Dog!
. Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.
The Ionic Bond Li + F Li+ F - 1 s 22 s 1 1 s 22 p 5 [He] 1 s 1 s 2[2 Ne] 2 s 22 p 6 e- + Li+ + Li Li+ + e- F F - Li+ F - 9. 2
The Ionic Bond Li + F Li+ F - 1 s 22 s 1 1 s 22 p 5 [He] 1 s 1 s 2[2 Ne] 2 s 22 p 6 e- + Li+ + Li Li+ + e- F F - Li+ F - 9. 2
Ionic Structures �In an ionic compound (solid), the ions are packed together into a repeating array called a crystal lattice. �The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other. �Its called simple cubic packing (Na. Cl is an example) �Ionic formulas are always Empirical Formulas (simplest)
Properties of Ionic Compounds • All ionic compounds form crystals. • Ionic compounds tend to have high melting and boiling points. To break the positive and negative charges apart, it takes a huge amount of energy. • Ionic compounds are very hard and very brittle. Again, this is because of the way that they're held together – strong attraction of oppositely charged ions. These ions simply don't move around - so they don't bend at all. This also explains the brittleness of ionic compounds. If we give a big crystal a strong enough whack with a hammer, we usually end up using so much energy to break the crystal that the crystal doesn't break in just one spot, but in a whole bunch of places. Instead of a clean break, it shatters.
• Ionic Compounds are poor conductors of electricity in the solid state. Ions are held tightly together and cannot move. Therefore ions cannot conduct electricity • Ionic compounds conduct electricity when molten (melted). Ions are mobile and can therefore conduct electricity. • Ionic compounds conduct electricity when aqueous (dissolved in water). If we take a salt and dissolve it in water, the water molecules pull the positive and negative ions apart from each other. Instead of the ions being right next to each other, they are able to move around in the water and conduct electricity. View here
Covalent Bonding
COVALENT BOND bond formed by the sharing of electrons
Covalent Bond �Between nonmetallic elements of similar electronegativity. �Formed by sharing electron pairs �Stable non-ionizing particles, they are not conductors at any state �Examples; O 2, C 2 H 6, H 2 O, Si. C
Covalent bonds- Two electrons. atoms share one or more pairs of outer-shell Fluorine Atom Fluorine Molecule (F 2)
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? F + 7 e- F F 7 e- 8 e- lone pairs single covalent bond lone pairs F lone pairs Dot diagram of F 2 F F lone pairs single covalent bond Lewis structure of F 2 9. 4
Lewis structure of water H + O + H single covalent bonds H O H H O or H 2 e-8 e-2 e. Double bond – two atoms share two pairs of electrons O C O or O O C double bonds - 8 e 8 e- 8 ebonds double Triple bond – two atoms share three pairs of electrons N N triple bond 8 e-8 e or N N triple bond 9. 4
Properties of Covalent Compounds • Covalent compounds tend to have low melting and boiling points. Many simple molecular substances are gases or liquids at room temperature. • Most molecular compounds are poor conductors of electricity in all states. There are no free electrons available to move and conduct an electric current. • Covalent compounds are usually much softer than ionic material. • Covalent compounds tend to be flammable than ionic compounds. • Most molecular compounds are insoluble in water, but will dissolve in nonpolar organic solvents.
9. 4
Drawing Lewis Structures 1. Write the dot diagram for each atom present in the compound. 2. Take the TWO atoms with the MOST unpaired dots and join (bond) them together. 3. Add the remaining atoms where they are needed the most. (starting with remaining atoms with most unpaired dots) Try the following: HOF N 2 H 4 CH 2 O H 4 CO HCN
Bonds in all the polyatomic ions and diatomics are all covalent bonds
In covalent bonding, one or more pair of electrons are shared. However, all ‘sharing’ is NOT the same. Therefore, we can have Nonpolar (Pure) Covalent Bonds and Polar Covalent Bonds.
NONPOLAR COVALENT BONDS when electrons are shared equally H 2 or Cl 2
Nonpolar Covalent bonds- Two atoms equally share one or more pairs of outer-shell electrons. Fluorine Atom Fluorine Molecule
POLAR COVALENT BONDS when electrons are shared but shared unequally. H 2 O
Polar Covalent Bonds: Unevenly matched, but willing to share.
- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms (electrons are shared unequally) electron poor region H electron rich region F e- poor e- rich H F d+ d-
To determine the type of Bonding present in a compound, we compare the electonegativities of the bonding atoms. The difference in electronegativity between these atoms predicts the type of bond present
The Electronegativities of Common Elements
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, F is highest electron poor region H electron rich region F
Classification of bonds by difference in electronegativity Difference Bond Type 0 0. 4 Nonpolar (Pure) Covalent > 1. 7 0. 4 < and ≤ 1. 7 Ionic Polar Covalent Creates a “Bonding Continuum” Increasing difference in electronegativity Nonpolar (Pure)Covalent Polar Covalent share e- equally partial transfer of e. Unequal sharing Ionic transfer e-
Classify the following bonds as ionic, polar covalent, or covalent: The bond in Cs. Cl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0. 7 Cl – 3. 0 – 0. 7 = 2. 3 H – 2. 1 S – 2. 5 – 2. 1 = 0. 4 N – 3. 0 – 3. 0 = 0 Ionic Nonpolar Covalent Non. Polar Covalent 9. 5
Predicting Molecular Geometry 1. Draw Lewis structure for molecule. 2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom. 3. Use VSEPR to predict the geometry of the molecule.
Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB 2 2 0 linear B B
0 lone pairs on central atom Cl Be Cl 2 atoms bonded to central atom 10. 1
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB 2 2 0 linear 0 trigonal planar AB 3 3
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry AB 2 2 0 linear trigonal planar tetrahedral AB 3 3 0 trigonal planar AB 4 4 0 tetrahedral
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom AB 3 3 0 AB 2 E 2 1 Arrangement of electron pairs Molecular Geometry trigonal planar bent
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom AB 4 4 0 tetrahedral AB 3 E 3 1 tetrahedral trigonal pyramidal Arrangement of electron pairs Molecular Geometry
VSEPR Class # of atoms bonded to central atom # lone pairs on central atom AB 4 4 0 Arrangement of electron pairs Molecular Geometry tetrahedral AB 3 E 3 1 tetrahedral trigonal pyramidal AB 2 E 2 2 2 tetrahedral bent O H H
bonding-pair vs. bonding pair repulsion < lone-pair vs. bonding < lone-pair vs. lone pair repulsion
Polarity and shape �The shape of the molecule directly influences the overall polarity of the molecule. �If there is symmetry the charges cancel each other out, making the molecule non-polar �If there is no symmetry, then its polar
�Polar bonds do not guarantee a polar molecule �Ex: CCl 4 and CO 2 both have polar bonds, but both are NON-POLAR molecules. They have a dipole moment of zero (no net dipole or ‘resultant vector’. �The greater the dipole moment, the more polar the molecule
Why is molecular polarity important? �Polar molecules have higher melting and boiling points (for example the BP of HF is 19. 5° C, and the BP of F 2 is – 188° C). �Polar solvents dissolve ionic and polar molecules more efficiently than non-polar solvents
Dipole Moments and Polar Molecules electron poor region electron rich region H F d+ d- 10. 2
‘polar molecule’
Which of the following molecules have a dipole moment? H 2 O, CO 2, SO 2, and CH 4 O H H dipole moment polar molecule S O O dipole moment polar molecule H O C O no dipole moment nonpolar molecule H C H H no dipole moment nonpolar molecule
Does BF 3 have a dipole moment? Symmetrical Molecule – No Resultant Dipole NONPOLAR MOLECULE
The bent shape creates an overall positive end and negative end of the molecule = POLAR The symetry of the molecule Cancels out the “charges” Making this NON-POLAR No overall DIPOLE
Examples to Try �Determine whether the following molecules will be polar or non-polar molecules. �SI 2 �CH 3 F �As. I 3 �H 2 O 2
Summary of Polarity of Molecules �Linear: �When two atoms attached to central atom are the same, the molecule will be Non-Polar (CO 2) �When the two atoms are different the dipoles will not cancel, and the molecule will be Polar (HCN) �Bent: �The dipoles created from this molecule will not cancel creating a net dipole moment and the molecule will be Polar (H 2 O)
Summary of Polarity of Molecules �Pyramidal: �The dipoles created from this molecule will not cancel creating a net dipole and the molecule will be Polar (NH 3) �Trigonal Planar: �When the three atoms attached to central atom are the same, the molecule will be Non-Polar (BF 3) �When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the molecule will be Polar (CH 2 O)
Tetrahedral �When the four atoms attached to the central atom are the same the molecule will be Non. Polar �When three atoms are the same, and one is different, the dipoles will not cancel, and the molecule will be Polar
METALLIC BOND bond found in metals; holds metal atoms together very strongly
Metallic Bond �Formed between atoms of metallic elements �Electron cloud around atoms �Good conductors at all states, lustrous, very high melting points �Examples; Na, Fe, Al, Au, Co
Metallic Bonds: Mellow dogs with plenty of bones to go around.
Ionic Bond, A Sea of Electrons
This ‘sea of mobile electrons’ allows metals to conduct an electricity current – electrons are free to travel. ‘Sea of Electrons’ is also the reason why most metals are malleable and ductile. As the electrons are free to move, it allows the metal to bend without breaking.
- Slides: 71