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- Chapter 3: Atomic Structure and Periodic Table
- 3.6 Electron Arrangements Within Atoms
Chemistry at a Glance: Shell-Sub-shell-Orbital
Interrelationships
- 3.7 Electron Configurations and Orbital Diagrams
3.8 The Electronic Basis for the Periodic Law and the
Periodic Table
- 3.9 Classification of the Elements
Chemistry at a Glance: Element Classification Schemes and the
Periodic Table
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- All s orbitals have the shape of a sphere, with its center at the
nucleus
- of the s orbitals, a 1s orbital is the smallest, a 2s orbital is
larger, and a 3s orbital is larger still
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- s orbital - a spherical-shaped atomic orbital; can hold a maximum of 2
electrons
- p orbital - a dumbbell-shaped atomic orbital; the three p orbitals (px,
py, pz) can hold a maximum of 2 electrons each
- Electrons always fill starting with the lowest-energy orbital:
- lower energy higher energy
- 1s2 2s2 2p6 3s2 3p6
- We will be concerned with only the valence electrons which are the
outermost electrons involved in forming bonds.
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- A p orbital consists of two lobes arranged in a straight line with the
center at the nucleus
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- Ground state electronic configuration of atoms in core format
- Carbon (C): ): [He] 2s2, 2p2
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or [He] 2s2, 2px13py13pz0
- Potassium (K): Ar] 4s1
- Phosphorous (P): [Ne] 3s2, 3p3
- Valence shell electronic configuration
- Carbon (C): ): 3s2, 3p2
- Potassium (K): 4s1
- Phosphorous (P): 3s2, 3p3
- How you get the electronic configuration of an atom from the periodic
table?
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- Carbon (C):
- Ground state: 2s2, 2p2
or 2s2, 2px13py13pz0
- Excited State: 2s1, 2p3 or 2s1, 2px13py13pz1
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- "octet rule“
- atoms tend to gain, lose or share electrons so as to have eight
electrons in their outer electron shell
- “Lewis structure of atoms”
- Shows only valence electrons, is
a convenient way of representing atoms to show their chemical bonding
pattern.
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- Covalent bonds - results from the sharing of electrons between two atoms
typically involves two nonmetallic elements
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- 3) Draw the skeletal structure by connecting the atoms with single
bonds.
- 4) Give each of the atoms an octet (8 e-). Adding unshared pairs of electrons
- 5) Count the total number of e- used through step 4 and compare to the
number calculated in step 2.
- a) If it results in zero, the
structure is correct.
- b) For every two electrons too
many, another bond is added
(minimize formal charges).
- Multiple bonds form only
with C, N, O and S.
- Total number of
bonds to neutral atoms:
- 4 bonds to C
- 3 bonds to N, P
- 2 bonds to O, S
- 1 bond to H, F, Cl, Br,
I
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- 1) Predict arrangement of atoms.
- H is always a terminal atom.
- Halogens and oxygen are often terminal.
- The central atom of binary compounds is usually written
- first and has the lowest subscript.
- Most organic compounds have more than two central atoms.
- These are mainly C, but N, O and S can also be central atoms.
- 2) Total number of valence electrons (e-)
- Add all valence electron of atoms in the molecule from the formula.
- Add the ion charge for negative ions or subtract for positive ions.
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- 1. For a neutral molecule, the sum of the formal charges equals
zero. For a polyatomic ion, the
sum of the formal charges equals the charge on the ion.
- 2. Formal charge of each atom is calculated by:
- (group #) - (# unshared e-)
- ½ (# shared e-)
- 3. Formal charges are shown as +
or - on the atom with that charge.
- 4. An atom with the same number of bonds as its group number has no
formal charge.
- 5. In a molecule if two different elements can be assigned a negative
charge, then the more electronegative element gets the charge; the same
sign should not be given to bonded atoms.
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- CHCl3
- C2H4
- C3H8O
- CH3CH2CH2OH
- CH3CH2OCH3
- CH3CO2H
- CH3CHO
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- For predicting shapes of molecules and polyatomic Ions based on the
repulsion of valence pairs of electrons making them as far apart as
possible around an atom of a
Lewis structure.
- 1) Draw the Lewis structure for the molecule or ion.
- 2) Determine the number of bonding and unshared pairs attached to the
central atom.
- One single, double or triple bond counted as a bonding pair
- 3) Choose the appropriate case from the given chart.
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- Hybridization is the mixing up of two or more atomic orbitals
- There are three types of hybrid atomic orbitals for carbon
- sp3 (one
s orbital + three p orbitals give four sp3 orbitals)
- sp2 (one s
orbital + two p orbitals give three sp2 orbitals)
- sp (one s orbital + one p
orbital give two sp orbitals)
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- H
- 2.1
- Li Be B C N O F
- 1.0 1.5 2.0 2.5 3.0 3.5 4.0
- Na Mg Al Si P S Cl
- 0.9 1.2 1.5 1.8 2.1 2.5 3.0
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- We classify chemical bonds as polar covalent, nonpolar covalent and ionic
based on the difference in electronegativity between the atoms
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- Identify covalent and ionic
compounds:
- NaCl, C2H5OH, CH3COOH, Na2CO3,
CH3OK, KOH
- Covalent :
- Ionic:
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- N-H nonpolar-covalent, polar-covalent or ionic bonds
- O-H nonpolar-covalent, polar-covalent or ionic bonds
- C-H nonpolar-covalent, polar-covalent or ionic bonds
- C-F nonpolar-covalent, polar-covalent or ionic bonds
- Na-Cl nonpolar-covalent, polar-covalent or ionic bonds
- Al-Cl nonpolar-covalent,
polar-covalent or ionic bonds
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- single bond - one shared pair of electrons between two atoms; a s bond
- double bond - two shared pairs of electrons between two atoms; one s
bond and one p bond
- triple bond - three shared pairs of electrons between two atoms; one s
bond and two p bonds
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- Count sigma bonds and unshared electrons around the atom
- If the total number of pairs:
- 2 sp hybridization
- 3 sp2 hybridization
- 4 sp3 hybridization
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- There are many compounds when dissolved in water changes the Hydrogen
ion (H+) (related to pH) concentration of to water to acid or
basic sides
- Binary acids: E.g. HF, HCl, HBr, HI, H2S
- Oxyacid: E.g. HNO3, H2SO4 , HClO4
- Organic acids: E.g.
- d) Hydroxy bases: NaOH, Ca(OH)2
- e) Amine bases: CH3NH2
(methylamine)
- (CH3)2NH
(dimethylamine)
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- This is the first acid/base concept to be developed to describe typical
acid/base reactions.
- Arrhenius Acid:
- A substance that produces H+,
or (protons) H+3O, (hydronium ion) in an aqueous
solution.
- Arrhenius Base:
- A substance that produces OH-,
or hydroxide ion in an aqueous
solution.
- E.g. HCl (acid), NaOH (base).
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- This is the second acid/base concept to be developed to include proton, H+
transfer reactions to base other than those containing OH-. This definition also uses
- conjugate acid/base concept .
- Bronsted Acid:
- A substance that donates
protons (H+): E.g. HCl (acid),
- Bronsted Base:
- A substance that accepts
protons. E.g. NH3
(base) non-hydroxy bases such as amines.
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- Lewis was successful in including acid and bases that catalyze organic
reactions without proton or hydroxyl ions.
- Lewis Acid: A substance that accepts an electron pair.
- Lewis base: A substance that donates an electron pair.
- E.g. BF3(g) + :NH3(g) ® F3B:NH3(s)
- Lewis Acid Lewis base Lewis acid/base adduct
- the Lewis base donates a pair of electrons to the acid forming a coordinate
covalent bond common to coordination compounds. Lewis acids/bases will
be discussed later in describing reaction mechanism
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- H2O(l) + HCl(aq) H3+O(aq) + Cl¯(aq)
- base acid conjugate
acid conjugate base
- NH3(aq) + H2O(l) NH4+ +
OH¯(aq)
- base acid
conjugate acid
conjugate base
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- Higher the Ka : higher
the acidity.
- Higher the Kb : higher
the basity.
- Higher the pKa : lower
the acidity.
- Higher the pKb : lower
the basity.
- Which one is weaker acid?
- HNO2 ; Ka= 4.0 x 10-4. pKa= 3.39
- HOCl2 ; Ka= 1.2 x 10-2. pKa= 1.92
- HOCl ; Ka= 3.5 x 10-8. pKa= 7.46
- HCN ; Ka= 4.9 x 10-10. pKa= 9.31
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- Equilibrium favors reaction of the stronger acid and stronger base to
give the weaker acid and the weaker base
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Notes
