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```									Practice

How many significant figures in each of the following?
138.20
0.6040
1.300 x 10-4
1050
1.06 x 103
How many significant figures are allowed in each of the
answers to the following arithmetic problems:
126.23 x 0.10
0.6831 + 125.26
3.632 / 0.0061
28 + 0.75
1
Last Time

Use of significant figures
Units and Measurement

Today
Dimensional analysis
using units to guide calculations

Atoms and Subatomic Particles, Chapter 5

2
Units / Dimensions

 Equivalence Statement
-relates one unit to another
1.00 metre = 100.0 cm        1 inch = 2.54 cm

 Conversion Factors
-relates one unit to another as a ratio for use in
mathematical operation
How many cms are in 3.75 metres?
100.0 cm        or     1.00 metres
1.0 metres                100.0cm

3.75 metres x 100.0 cm  375cm
1.00 m                                  3
Dimensional Analysis

 A problem solving technique using the units as a guide
-use the units in calculations
-identify the given information, number and units
-identify the units required for the correct answer
-set up equivalence statements and or conversion factors
-do the math, cancel out units where possible

If the solution has been set up correctly, the correct units
should be generated in the process.

4
Physical Properties as Conversion Factors

 Density       mass             g     g      kg
unit volume        mL    cm3     L

Density is specific for temperature, that is density
changes with temperature. Standard is 25 oC

 Specific Gravity

density of a substance     g/mL
unitless
density of water         g/mL

5
Density

 A copper (Cu) coin has a mass of 3.14 g. The density of copper is
8.96 g/cm3. What is the volume of the coin?
given quantities and units:
mass 3.14 g   density 8.96 g/cm3

quantity and units required:     volume: cm3

equivalence statement:    1 cm3 = 8.96 g

set up conversion factor: 1 cm3 / 8.96 g     or     8.96 g / 1 cm3

calculation and cancel units:                 3

3.14 g x 1 cm  0.350 cm
3

8.96 g

significant figures: 3 sig figs allowed in answer
6
Example:
Percent as a Conversion Factor

 How many grams of water are contained in 65.3 g of a mixture of water and
alcohol that is 34.2% H2O by mass?

Given:         65.3 g of mixture , 34.2 % water ; 100 g mix = 34.2 g water

Required:      grams of water

Conversion factor:      34.2 g water         100g mixture
100 g mixture         34.2 g water
Calculation:       34.2 g water
x 65.3 g of mixture
100 g mixture

Significant figures: 100 comes from 100%, considered an exact
number. Answer allowed 3 sig figs
7
Atoms and Subatomic Particles
Chapter 5

 Element: a pure substance that cannot be broken down
into simpler pure substances

 Compound: a pure substance that can be broken down to
two or more simpler substances

 Atom: the smallest particle of an element that still has the
properties of the element

8
John Dalton         19th century

Experimental observations:
-most natural materials are mixtures of pure substances

-pure substances are either elements or compounds

-a given compound always contains the same proportion,
by mass of the elements
(The law of definite proportions)

e.g.:
water always contains 8 grams of oxygen for every 1 gram
of hydrogen

9
Dalton’s Atomic Theory               (original early 1800’s) (Atomic
Theory of Matter) pg. 123

 Modern Version
 All matter is made up of small, neutral particles called atoms. There
are 112 different types of atoms. Each type corresponds to a different
element.
 All atoms of a given element are similar to one another.
 Atoms of a given element are different from those of any other
element.
 Atoms of one element combine with atoms of other elements to form
compounds.
 A compound always has the same relative numbers and types of atoms.
 During a chemical reaction, changes occur only in the way atoms are
grouped together
ie: atoms are neither created nor destroyed in chemical reactions

10
Prediction:

Atoms of a given pair of elements could combine in different
proportions and produce different compounds.

e.g:   N nitrogen and O oxygen
1:1     NO    nitric oxide, colourless gas

1:2     NO2 nitrogen dioxide, brown gas

2:1     N2O nitrous oxide, colourless gas, (laughing gas)

3 different pure substances (compounds) each with its own
distinct set of chemical and physical properties but made up from the
same two elements.
11
Charge

 A piece of amber rubbed with cloth will attract some
objects to it and repel others.
 A glass rod rubbed with silk will attract some objects and
repel others.
 The glass rod will repel other glass rods that have been
rubbed with silk; the amber will repel other pieces of
amber that have been rubbed with cloth. The glass rod and
the amber will attract each other.
 The amber and the glass rod have become “charged”
 By convention the amber is assigned a NEGATIVE
charge.
 The glass rod has a POSITIVE charge
 Like charges repel, opposite charges attract
12
Structure of the Atom -Subatomic Particles

Electrons:       1890’s J. J. Thomson

observations:
-atoms of any element could be made to emit tiny negative
particles
-same (-)ve particle no matter which element

conclusion:
-all atoms contain negative particles (called electrons),
that have almost negligible mass

13
J.J. Thomson (1856-1940):

» experimented with Cathode-ray tubes
» beam could be deflected with an applied electrical field or
applied magnetic field
» deflected in manner expected for negatively charged particles

14
Last Time

 Dalton’s Atomic Theory
 Subatomic Particles
– JJ Thompson, cathode rays, electrons

Today
 other subatomic particles
 modern view of the atom
 describing an atom

15
J.J. Thompson

 Conclusions:
– cathode rays are negatively charged fundamental
particles of matter found in all atoms
– all atoms contain these negative particles (electrons)
– measured charge-to-mass ratio of an electron

– http://www.howstuffworks.com/tv2.htm

16
J.J’s Atomic Model

-electrons distributed randomly in a diffuse positive
cloud.
-“plum pudding” model: raisins dispersed in pudding.

Plum Pudding model
17
Subatomic Particles

Nucleus:      1911 Ernest Rutherford
-atoms are neutral therefore there must be a positive
component

experiment:
-directed positive alpha particles (heavy particles with a
+2 charge) toward a thin metal foil

prediction:
-large, fast alpha particles would pass straight through with
no deflection

18
Rutherford’s gold foil experiment
02_24

Some alpha            Most particles
particles are         pass straight
Uranium source of                              scattered             through foil
alpha particles (embedded
in a lead block to absorb

Thin
Beam of
metal foil
alpha particles   Luminescent screen
to detect scattered
alpha particles
 particles

19
Subatomic Particles

 observations:
-most alpha particles passed straight through
-some were deflected with large angles
-some were reflected straight back

 conclusion:
-there must be a large centre of concentrated positive
charge in the atom

20
Subatomic Particles: The Nucleus

 Nucleus:
-dense centre of positive charge surrounded by moving
(-)ve electrons.
-positive charge must balance the negative charge of the
total number of electrons

 Protons:
-particle of (+)ve charge, contained in the nucleus
-same magnitude of charge as the electron (e-)
-much greater mass than an electron (~1800 x mass of e-)
-number of protons must be equal to the number of
electrons
21
Subatomic Particles

-most nuclei also contain neutral particles called neutrons

-slightly larger mass than a proton but no charge

-neutrons and protons collectively called nucleons

- some elements can have atoms with different numbers
of neutrons

22
Revise Dalton’s Atomic Theory

All atoms of the same element contain the same number of
protons and electrons.

but

Atoms of a given element may have different numbers of
neutrons.

All atoms of a given element are similar to one another

23
Modern View of the Atom

The Nuclear Atom
 nucleus: -dense centre of positive charge
-contains (+)ve charged particles, protons,
and neutral particles, neutrons
 electrons: -particles of (-)ve charge, fill the space around
the nucleus

 mass: - protons and neutrons make up most of the mass of
an atom

 neutral: - same number of protons and electrons

The chemistry of an atom arises from its electrons          24
Description of an Atom

 Atomic Number
-characteristic of the element
-equal to the number of protons in the nucleus
(therefore also equal to the number of electrons in the
neutral atom

 Mass Number
-equal to the number of protons + the number of
neutrons in the nucleus of the atom

25
Representation of an Element

A
Z   X

Where X is the symbol of the element
A is the mass number (always a whole number)
Z is the atomic number

12        14           16
6   C     7   N        8   O

so this carbon has 6 protons, 6 electrons and 6 neutrons
26
Isotopes
-atoms of the same elements with the same number of
protons, but with different number of neutrons
12        14        1                 2     3
6   C,    6   C    1   H,         H, 1     1H
D, deuturium   T, tritium

Isobars
-atoms that have the same mass number but different
atomic number (so they are different elements)
14        14
6   C,    7   N

27
Mass of an Atom

 Atomic Mass (weight)
eg: one atom of 12C weighs 1.99 x 10-23 grams

 Atomic Mass Unit – AMU
a unit of mass, set relative to a standard mass

definition: mass of 1 atom of   12C   = 12.00000..amu

12.0000 amu = 1.99 x 10-23 grams
1 amu = 1.66 x 10-24 grams

28
Atomic Masses of the Elements

 atomic masses of the individual elements are generally
shown on the periodic table

 represents the average mass, in amu, of an atom of the
element, considering the relative abundance of the isotopes

Atomic mass of C given as 12.01 amu
12C 98.98%          12.000 amu            11.866
13C 1.11%           13.003 amu             0.144
14C <0.01 %         14.003 amu             0.001
Weighted average, atomic mass of C: 12.011 amu

29
Other Atomic Species

 Ions
-obtained when electrons are added to, or removed from a
neutral species

 Cation
-a positively charged ion
-the result of removing 1 or more electrons from a neutral
species
Li  Li+ + e-
 Anion
-a negatively charged ion
-the result of a neutral species gaining one or more electron
O + 2e-  O2-                          30
Ions

Simple ions:
– charged species containing only one atom:
eg:    Cl- , Na+ , Ca2+, O2-

Polyatomic ions:
– charged species containing a group of atoms held
together by strong bonds:

NO2-, CN- , NH4+, SO42-

see Table 8.4 Chapter 8 page 263

31
 The formation of ions is only ever due to the addition or
loss of electrons.

 Protons cannot be removed or added to atomic or other
species under normal chemical circumstances

32
Structural Units of Pure Substances

 the smallest units:

atoms      Add or remove electrons    ions
(neutral)                        (charged species)

These small units combine in various ways to produce pure
substances

33
Structural Units of Pure Substances

1) Combinations of Neutral Atoms
atoms are the building blocks of matter
a) Elements
-contain only one kind of atom, grouped together
-an extended array of individual atoms
Atomic solids
eg: gold, carbon , iron

Atomic gases
eg: argon , neon

Atomic liquid
eg: mercury
34
Structural Units of Pure Substances

b) Molecules:
-group of two or more atoms that function as a unit
(molecule)
-atoms are tightly bound together in the unit
-each molecule behaves as an individual unit or particle
Elements
-all atoms are the same, atoms are grouped in
molecules
eg: O2 oxygen, N2 nitrogen, Cl2 chlorine
Compounds
-more than one type of atom in the molecule
H2O, CO2 ,
35
Structural Units of Pure Substances

2) Combination of ions:

Ionic Compounds
-from combinations of anions and cations to
form neutral species

eg: Na+ and     Cl- give NaCl

Mg2+ and Cl- give MgCl2

36
Natural States of Elements

 Most elements are very reactive and so tend to exist as
compounds
 Some are found in pure form in nature

 Noble Metals
-inert, that is unreactive, very stable
- eg: gold, silver, platinum, palladium

Noble (Inert)Gases
-eg: helium, neon, argon, krypton, xenon

Air
-oxygen, nitrogen, hydrogen
37
Writing Chemical Formula

Chemical Formula
-chemical representation of a substance
-use chemical symbols of the atoms
-gives some indication of the relative numbers and types of
atoms or ions in that substance

38
Chemical Formula
 Elements (extended array of atoms)
-represented by the element symbol
Ag, Au, Ne, Fe

 Molecular compounds
-element symbols plus numerical subscript giving the
number of atoms of each element present in one
molecular unit
Elements: molecular units: H2 O2 N2 F2 Cl2 Br2 I2
(P4) (S8)
Compounds: eg: H2O, CCl4,, P2O4
-formula does not give any information about how the
atoms are arranged or joined together, or anything
about the shape of the molecule                        39
Chemical Formula

 Ionic Compounds
-represent the simplest whole number ratio of anions to
cations
charge must balance to give a neutral species
Combination of simple ions:
NaCl , MgCl2 ,
Combination of simple ions and polyatomic ions:
Na+ and SO42- Na2SO4 ,
Ca2+ and PO43- Ca3(PO4)2
Combination of polyatomic anions and cations
NH4+ and NO3-       NH4NO3

40
Summary

Matter
physical techniques

pure substances
elements                                    compounds
chemical means

Aggregate of   Homoatomic                     Molecular        Ionic
like atoms     molecules                    compounds      compounds
Cu                O2                         H2O           NaCl
Au                F2                         CO2           CaF2
Ag                N2                         NO2           (NH4)2CO3

41

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