Atomic theory J' Dalton, 1803

1
Atomic theory (J. Dalton, 1803)

• 1) all matter is composed of small indivisible
  particles called atoms
• 2) In chemical reactions the atoms of one
  element cannot be changed into atoms of another
  element
• 3) atoms can be neither created nor destroyed
• (basis for the Law of Conservation of Matter)
• This happens in nuclear reactions (e.g.
  reactors, radioactive decay, stars, bombs E
  mc2)

2
Atomic theory (J. Dalton, 1803)

• 4) atoms of the same element are identical in
  mass, size and properties (except for isotopes)
• 5) atoms of one element differ in mass and other
  properties from atoms of other elements
• 6) atoms can combine chemically to form
  compounds, which contain atoms of different
  elements in whole number ratios (e.g. H2O
  C6H12O6)
• (law of constant composition)

3
Atomic theory (J. Dalton, 1803)

• 7) Chemical reactions result in different
  combinations of the atoms in the starting
  materials to yield new products.
• C8H18 O2 CO2 H2O
• ( a chemical equation)
• 2 C8H18 25 O2 16 CO2 18 H2O
• ( a BALANCED chemical equation)

4
Modern Atomic Theory view of matter
Modern Atomic Theory view of a chemical
reaction
5
The Nuclear Atom Model

• atoms consist of subatomic particles that are the
  same for all atoms electrons (e-), protons (p)
  and neutrons (n).
• The nucleus is DENSE (contains p and n) and is
  surrounded by a diffuse electron cloud
• different elements have different atomic
  compositions (e.g. H atoms vs. He atoms)

6
The Nuclear Atom Model

• mass of e- 9.11 x 10-28 g (Table 2.1 p 48)

the charge/mass ratio of p was determined by
mass spectroscopy assuming same charge as that
on e- the mass of a proton 1.672 x 10-24 g
(1830 x greater than for e-)
7
The Nuclear Atom Model

• 1911 Rutherford proposes nuclear model for atom
  dense () charged core surrounded by diffuse (-)
  charged electrons, but only a portion of the
  nuclear mass was accounted forproposed existence
  of neutron
• 1932 neutron discovered by Chadwick (mass
• of neutron 1.675 x 10-24 g )

8
See Figure 2.8 Rutherfords Experiment
9
Summary of the Atom

• atoms are the smallest particles that can be
  uniquely associated with an element
• each element has unique atoms (for isotopes see
  below)
• atoms are composed of e-, p and n
• atoms are electrically neutral ( of e- of p)
• for a single element, isotopes differ only in
  number of n (neutrons)
• atoms have characteristic masses (atomic weights)
• atoms combine with one another in definite, whole
  number proportions to make compounds

10
The Nuclear Model of the atom
Table 2.1
11
Notation for Atoms

• 12C
• 13C
• C

only one isotope of carbon
only one isotope of carbon
all isotopes of carbon
12
12C
Figure 2.9 Isotopes of carbon
13C
13
Using the Periodic Table

• atomic number number of protons in the
  nucleus corresponds to the position of that
  element in the periodic table
• atomic weight average mass of an atom
  calculated from the masses and natural abundances
  of all isotopes
• (use atomic weights to calculate the molecular
  weights of compounds from their constituent
  elements!)
• mass number sum of protons neutrons in the
  nucleus
• isotopic mass mass of a single isotope

14
Atomic weight measurements

• How was the atomic weight measured?
• By mass spectrometry
• This also measures
• natural abundance
• for a given isotope

Figure 2.13 Mass spectrum of Ne
15
Atomic weight calculation (p 52)

• There are three naturally occuring isotopes of
  neon (Ne)
• 20Ne mass 19.99244018 amu
• 21Ne mass 20.9938467 amu
• 22Ne mass 21.9913855 amu
• the atomic weight is reported in text as
• 20.1797 amu

16
Atomic weight calculation

• How was the atomic weight calculated?
• multiply each isotopic mass by the reported
  natural abundance for the isotope, then
• add these individual contributions for each
  isotope to get the average atomic weight for the
  element

17
Atomic weight calculation

• There are three naturally occuring isotopes of
  neon (Ne)
• 20Ne mass 19.99244018 amu (90.51)
• 21Ne mass 20.9938467 amu (0.27)
• 22Ne mass 21.9913855 amu (9.22)
• the atomic weight is reported in text as
• 20.1797 amu
• 18.09515 0.05668 2.0276 20.1794 amu
• 20.18 amu

18
Atomic Masses

• 13C
• 12C

13.00335 amu (1.11)
12.0000 amu (98.89)
atomic weight of C 12.01 amu WHY? (more
accurate average mass of C 12.01115 amu)
19
Atom Summary (know for Exams)

• atoms are the smallest particles that can be
  uniquely associated with an element
• each element has unique atoms (for isotopes see
  below)
• atoms are composed of e-, p and n
• atoms are electrically neutral ( of e- of p)
• for a single element, isotopes differ only in
  number of n (neutrons)
• atoms have characteristic masses (atomic weights)
• atoms combine with one another in definite, whole
  number proportions to make compounds

20
Chemical Reactions
Figure 2.28
reactants products reactants
starting material products products of the
reaction
21
Chemical Formulas

• How many atoms of oxygen in this compound?
• CaCO3
• How many atoms of oxygen in this compound?
• Al2(SO4)3

22
Chemical Formulas

• Al2(SO4)3 is equivalent to Al2S3O12
• the term (SO4)3 in the chemical formula indicates
  the presence of three sulfate ions in the
  compound.
• This notation indicates a specific grouping of
  the S and O atoms to form the sulfate ion.

23
Chemical Equations

• identify reactants and products (i.e. show
  chemical formulas for each species)
• show stoichiometric relationships between all
  chemical species in the reaction
• Chemical equations must be (mass) balanced!

24
Chemical Equations

• CH4 2O2 CO2 2H2O
• coefficients subscripts
• coefficients can be changed to achieve mass
  balance but subscripts are never changed to
  balance an equation, because a change in a
  subscript implies a change in chemical
  composition.

25
Balancing a Chemical Equation

• write formula for each reactant and product on
  the correct side of the reaction arrow
• counts atoms of each element on both sides of
  arrow
• start with the compound which has the most
  complex formula
• add coefficients to chemical formulas to balance
  numbers of each atom
• trial and error begins...

26
Summary of the Atom

• atoms are the smallest particles that can be
  uniquely associated with an element
• each element has unique atoms (for isotopes see
  below)
• atoms are composed of e-, p and n
• atoms are electrically neutral ( of e- of p)
• for a single element, isotopes differ only in
  number of n (neutrons)
• atoms have characteristic masses (atomic weights)
• atoms combine with one another in definite, whole
  number proportions to make compounds

27
Figure 2.18 Covalent compounds
28
Figure 2.21 Ionic compounds
29
(No Transcript)
30
Table 2.4
Within a main group, the common ions have
similar charges (e.g. 1, 2, 3, 2-, 1-, etc.)
31
A portion of Table 2.5
Some polyatomic ions.
32
2.6