Introduction: Matter, Measurement and Moleculescheminnerweb.ukzn.ac.za/Files/Chem 110 (2012)/Chem 110 2012 Ch1... · “study of matter & changes it undergoes ... Changes in matter

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Chapter 1

Introduction: Matter,

Measurement and Molecules

Dr V Paideya

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

CHEMISTRY

“study of matter & changes it undergoes”

“matter is anything that has mass and takes up space”

- study of physical & chemical properties of matter

- what changes occur in these properties, in the course of/as the result of a

chemical reaction, & how these changes may be observed

- why the reaction involved does (or doesn’t…) occur

be able to understand & explain such

macroscopic changes from an atomic/molecular

(submicroscopic) perspective

States (Phases) of Matter

- solid, H2O(s); liquid, H2O(l); gas, H2O(g)

- phase transitions occur @ specific P/T values,

governed by properties of atoms/molecules


Matter

- atoms are building blocks of matter

- each element is made of same kind of atom/molecules

- compounds made of two or more different kinds of elements bonded together


Pure Substances, Elements & Compounds

pure substance

-has distinct properties & unvarying/constant composition

eg. NaCl(s), H2O(l), HCN(g)

element

-substance that cannot be decomposed into simpler substances

eg. Cl2(g), Br2(l), I2(s); Ne(g), Hg(l), Au(s)

compound

-substance composed of 2 or more different elements

2 or more different kinds of atoms

eg. UF6(g), H2O(l), CaCO3(s)

Law of Constant Composition/Definite Proportions (Joseph Proust ca 1800)

“...elemental composition of pure substance is always the same…”

- different samples of pure compound have the same elemental

composition

- elements present in such samples have same proportion by mass


Mixtures

- combination of 2 or more substances, in which each substance retains own

chemical identity & can thus be separated from each other

- 2 types:

heterogeneous:

- mixture of visibly different composition, properties or appearance

eg. sand in H2O(l) (s, l), sand & NaCl (s, s), petrol & H2O(l) (l, l)

homogeneous:

- mixture of visibly uniform composition, properties & appearance throughout

eg. NaCl(aq) (s,l), air (g,g), stainless steel (s,s), soda water (g,l)

Properties:

- physical: measurement without changing identity/composition eg. mass, , v

-chemical: must involve change in chemical identity eg. flammability, reactivity

- extensive: dependent on quantity of sample involved eg. mass, volume

- intensive: independent of quantity eg. , m.p., b.p.; useful for identification


Classification of Matter

Figure 1.5


Changes of Matter

Physical Changes

Changes in matter that do not change the

composition of a substance.

Changes of state, temperature, volume, etc.

Chemical Changes

Changes that result in new substances.

Combustion, oxidation, decomposition, etc.


Chemical Reactions (Chemical Change)

In the course of a chemical reaction, the reacting

substances are converted to new substances.

Figure 1.6


Compounds

Compounds can be broken down into more

elemental particles; for example, during the

electrolysis of water, the smaller particles

hydrogen gas and oxygen gas are created.


Separation of Mixtures

1. Distillation

Separates a

homogeneous

mixture on the

basis of differences

in boiling point.

Figure 1.8


Separation of Mixtures

2. Filtration

3. Chromatography

Separates substances on the basis of differences in

solubility in a solvent

Separates solid substances from liquids and solutions.


The Scientific Method

A systematic approach to solving problems


SI Units

Système International d’Unités

Uses a different base unit for each quantity


Metric System

Prefixes convert the base units into units that

are appropriate for the item being measured.

Tera T 1012


SI Units - Temperature

The Kelvin is the SI

unit of temperature.

It is based on the

properties of gases.

There are no negative

Kelvin temperatures.

K = C + 273.15

Figure 1.10


Derived SI Units VolumeUnits of Measurement

Derived SI Units: Volume, V; E

-1 m10 d(eci)m 100 c(enti)m 1000 m(illi)

-1 L 1 dm3 (a cube 1 dm x 1 dm 1 x dm)

-1 mL 1 cm3(a cube 1 cm x 1 cm 1 x cm)

-1 kL 1 m3 (a cube 1 m x 1 m x 1 m)

ie. 1 m3 103 dm3 106 cm3 109 mm3

1 kL 103 L 106 mL 109 (micro)L

Derived SI Units: Density, or d; I

- physical property; units are usually

-g cm-3 (g/cm3, g mL-1, g/mL) for liquids

& solids

-g dm-3 (g/dm3, g L-1, g/L) for gases


Derived SI Units Density

Density is a physical property of a substance

and is determined through the following

formula:

density =mass

volume

=mVor symbolically


Uncertainty in Measurements

Different measuring devices have different

uses and different degrees of accuracy.

Figure 1.12


Significant Figures

All digits of a measured quantity,

including the uncertain, are called

significant figures.

The greater the number of significant

figures, the greater the certainty of the

measurement.


Significant Figures All nonzero digits are significant, e.g. 123.45

Zeros between two significant figures are

themselves significant, e.g. 103.405

Zeros at the beginning of a number are

never significant, e.g. 00123.45 = 123.45

Zeros at the end of a number are

significant if a decimal point is written in the

number, e.g. 123.450 has six significant figures but 123450 has

only five significant figures


Significant Figures

When addition or subtraction is

performed, answers are rounded to the

least significant decimal place.

When multiplication or division is

performed, answers are rounded to the

number of digits that corresponds to the

least number of significant figures in any

of the numbers used in the calculation.


Precision and Accuracy

Accuracy refers to the proximity of a measurement

to the true value of a quantity.

Precision refers to the proximity of several

measurements to each other.

Figure 1.15


Significant Figures1. any figure that is not zero is significant:

snot/units 845 mL _____ s.f.1243.29 mg _____ s.f.

2. zeroes between non-zero figures are significant:1906 mL _____ s.f.40501.09 J _____ s.f.

3. exact (“counting”) numbers by definition have an number of s.f., so physical

constants defined to be exact numbers do so also...:1 atm 101.325 kPa 760

mmHg; 0 OC 32 OF 273.15 K all _____ s.f.

4. leading zeroes (to the left of the first non-zero figure) are not significant:

snot/units 0.008 kg _____ s.f.0.003798 L _____ s.f.

5. trailing zeroes (to the right of the last non-zero figure) are significant only if the

number has a d.p.: 300.0 mm _____ s.f.0.0300 mm _____ s.f.

6. in measurements without a d.p., the number of s.f. is ambiguous:1200 mm ?? either: i) use snot

OR ii) 1200.


Using Significant Figures in Calculations

- all calculations governed by two fundamental rules

multiplication/division

- number of s.f. in final answer is the LEAST of numbers of s.f. in each of

original measurements

addition/subtraction

- number of d.p. in final answer is the LEAST of numbers of d.p. in each of

original measurements

Eg. 1

Calculate

i) volume, in mm3, of a box of length 6.741 cm, breadth 2.441 x 10-1 m, & height

4.2 mm i) 6.9 x 104 mm3;

ii) density () of a pure liquid, in g cm-3, if 103.67 g of it is needed to fill the box

completely ii) 1.5 g cm-3

Eg. 2

An empty container of mass 23.29 g has a mass of 86.1 g when filled with 0.5000

dm3 of a pure liquid. Determine the of this liquid in g cm-3. 0.126 g cm-3


Early Atomic Theory

John Dalton 1803 - 1807

-each element is composed of very small, indestructable, particles called atoms*

-all atoms* of given element are physically & chemically identical to each other, but

atoms of a particular element are different from atoms of all other elements

-Law of Conservation of Mass

-atoms are neither created or destroyed in chemical reactions

- mass reactants present @ start = mass products formed @ completion*

-Law of Constant Composition

-different samples of a pure compound have the same elemental composition

-elements present in such samples have same proportion by mass

-Law of Multiple Proportions

-if 2 elements (C & O) can combine to form 2 or more different compounds (CO &

CO2), the different masses of one element (O) combining with a fixed mass of the

other (C) can be expressed as a simple integral ratio


Atomic TheoryThe theory that atoms are the fundamental building blocks of matter came

into being during the period 1803 to 1807 in the work of John Dalton.Dalton’s Postulates Each element is composed of extremely small particles called atoms.

All atoms of a given element are identical to one another in mass and other properties, but the atoms of a particular element are different from the atoms of all other elements.

Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. This is the basis of the law of conservation of mass (or law of conservation of matter) which states that the total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. This is the basis of the law of constant composition (or law of definite proportions) which states that the relative numbers and kinds of atoms are constant, i.e. the elemental composition of a pure substance never varies.


The Law of Multiple Proportions

Was deduced by Dalton from the preceding four

postulates and states that:

If two elements A and B combine to form more

than one compound, the masses of B that can

combine with a given mass of A are in the ratio

of small whole numbers.

Examples

H2O consists of 2 hydrogens and 1 oxygen

H2O2 consists of 1 hydrogen and 1 oxygen


The Discovery of Atomic Structure

Cathode Rays & Particles (JJ Thomson, 1897)

- electrical discharges from cathode originally thought to be new form of radiation

- showed that radiation emitted was

- independent of cathode material used

- deflected by magnetic/electric fields

- findings consistent with model in which “beam” /”rays” composed of negatively

charged “particles”(-) with charge/mass ratio of - 1.7588 x 108 C g-1, or

-5.6857 x 10-9 g C-1

Electron Charge & Mass (Robert Millikan, 1909)

- oil drop experiment

- (-) charge on oil drops found always to be a multiple of

minimum value of 1.6(02) x 10-19 C ie. 1.602 x 10-19 C must

be charge of single electron (e-)

-mass of single e- determined to be 9.109 x 10-28 g

- only 1/1836 of mass of an H atom

- first subatomic particle


Radioactivity

The spontaneous emission of radiation by

an atom was first observed by Henri

Becquerel. It was also studied by Marie

and Pierre Curie.


Radioactivity

Three types of radiation were discovered by

Ernest Rutherford

particles

particles

rays

Figure 1.21


Discovery of the Nucleus

Ernest Rutherford shot particles at a thin sheet of gold foil and

observed the pattern of scatter of the particles.The Nuclear Atom

Some particles were deflected at large angles.

This led Rutherford to postulate that the

atom had a nucleus.


Modern Atomic Structure

- more than 99.99 % of atom mass & entire Q+ centred in atomic nucleus,

where nucleons (protons, p+ (Rutherford, 1919) & neutrons, nO (Chadwick,

1932) are collectively bound together by strong nuclear force

- atomic nucleus surrounded by much larger atomic volume, containing as

many e- as p+, so atom is electrically neutral & held together by force of

Coulombic/ electrostatic attraction


element symbol

(number of e- in neutral atom)

atomic number (Z)

number of p+

mass number (A)

number of p+ & nO

ZA E

Atomic (Z) & Mass (A) Numbers

- atoms of different elements have different numbers of p+ in their nuclei


Isotopes

Atoms with identical atomic numbers (Z) but different

mass numbers (A), or atoms with the same number of

protons which differ only in the number of neutrons are

called isotopes.

Examples:

116C

126C

136C

146C

carbon-12 isotope

carbon-14 isotope


Isotopes

- atoms of same element having different numbers of nO in their nuclei

- ie. same Z, different A, or same Z, different N

- chemical properties largely similar, but physical properties, & particularly the “nucular”

ones involving radioactive “nuculei”, can be very different

- each Mg ( ) atom is one of three naturally occurring isotopes

- 24Mg; 25Mg; 26Mg

Mg-24 Mg-25; Mg-26

-

-three isotopes of H individually named:

protium (1H, 1 p+/0 nO); deuterium (2H, 1 p+/1 nO); tritium (3H, 1 p+/2 nO)

- four isotopes of C:

11C; 12C; 13C (NMR probes & MRI scanners); 14C(“radiocarbon” dating)


Isotopes

Eg. 3 Complete the table below:

Experimentally..

- High Resolution Mass Spectrometry (p. 28) used for very precise (4-6 d.p.; 7-10

s.f. in total) measurements of the masses of an element’s isotopes & their naturally

occurring abundances

do: Q 1.44, 1.48 - 1.50

Element name Symbol p+ No e- A

79197Au

Ba 138

143 92

Pb 126

Krypton36


Atomic Mass

Atomic and molecular masses can be

measured with great accuracy with a mass

spectrometer.

Figure 1.23


Average Atomic Mass(commonly called Atomic Mass) We use average masses in calculations, because we use

large amounts of atoms and molecules in the real world.

Average atomic mass is calculated from the fractional

abundance of each isotope and mass of that isotope.

For example, the average atomic mass of C -

made up mostly of 12C (98.93%) and 13C (1.07%) - is 12.01

u.

extremely small SI masses of individual atoms (~4 x 10-22 g) too awkward

for everyday usage, so masses expressed in unified atomic mass units

(amu, u):

1 amu (u) = 1.66054 x 10-24 g 1 g = 6.02214 x 1023 amu (u)


Average Atomic Masses of Naturally Occurring Elements

-use average masses in “real world” calculations, as even smallest weighable sample

(~1 g 10-6 g) involves gobsmackingly large (~1015, or 10 quadrillion) numbers of

atoms

-no AAMs calculated as “weighted average” of an element’s isotopic masses (IMs) &

naturally occurring abundances

AAM= (IM x % ab/100) or (IM x fr ab)

Eg. 4

a)Naturally occurring Mg has three isotopes: 24Mg 78.99 %, 23.9850 u25Mg 10.00 %, 24.9858 u

Calculate its AAM. 26Mg 11.10 %, 25.9826 u

b) Naturally occurring Pb has four isotopes: 204Pb 1.40 %, 203.973037 206Pb 24.10 %,205.974455207Pb 22.10 %, 206.975885

Calculate its AAM. 208Pb 52.40 %, 207.976641


Eg. 5

Chlorine has two naturally occurring stable isotopes: 35Cl 34.968853 u 37Cl 36.965803 u

If the (average) atomic mass of naturally occurring elemental Cl is 35.453 u, what

are the % abundances of the two isotopes?

http://www.sisweb.com/referenc/source/exactmaa.htm

do: antimony, chromium, & nickel AM from IM’s & abundances

copper, rubidium abundances from IM’s & AM*

The Periodic Table

- rapid post-Dalton growth in experiment-based chemical knowledge showed very

quickly that many elements could be grouped together on basis of

similarities in their physical & chemical properties

- arrangement of elements in order of Z showed that these similarities occurred

in repetitive/periodic patterns, & agreed so closely with experimentally acquired

data, that phys/chem properties of 2 “missing” elements were accurately predicted

before their being reported as formally discovered, &/or phys/chem properties

characterized




Eg. 4

a) Naturally occurring Mg has three isotopes: 24Mg 78.90 %, 23.9850 u25Mg 10.00 %, 24.9858 u

Calculate its AM. 26Mg 11.10 %, 25.9826 u

AM = (IM x % ab/100)

= {[23.9850 x (78.90/100)] + [24.9858 x (10.00/100)] + [25.9826 x (11.10/100)] }

= 18.924 + 2.4986 + 2.8841

= 24.31 u (amu) [24.306]

b) Naturally occurring Pb has four:204Pb 1.40 %, 203.973037206Pb 24.10 %, 205.974455207Pb 22.10 %, 206.975885

Calculate its AM. 208Pb 52.40 %, 207.976641

AM = (IM x % ab/100)

= 2.856 + 49.640 + 45.742 + 108.98

= 207.2 u (amu)

[207.22]

back


Eg. 5

Chlorine has two naturally occurring stable isotopes:35Cl 34.968853 u 37Cl 36.965803 u

If the (average) atomic mass of naturally occurring elemental Cl is 35.453 u, what are

the % abundances of the two isotopes?

Assume that the fractional abundance of 35Cl is z......then the fr ab of 37C is

1-z ( fr ab = 1)

AM = (IM x fr ab)

35.453 =[(34.968853) x z] + [(36.965803) x (1-z)]

=34.968853 z + 36.965803 - 36.965803 z

-1.5128 = -1.996950 z (35.453 - 36.965803 -1.5128)

z =0.7576 3 dp 6 dp

% ab 35Cl =75.76%

% ab 37Cl =24.24 %


The Periodic Table

When one looks at the chemical properties of elements, one notices a

repeating pattern of reactivities.


Periodic Table The rows are called periods.

The columns are called groups.

Elements in the same group have similar chemical

properties.

Nonmetals are on the right side of the periodic table (with

the exception of H).

Metalloids border the stair-step line (with the exception of Al

and Po).

Metals are on the left side of the chart.


Groups

The above five groups are known by their names.

Table 1.7


The Periodic TableMetals, Non-Metals, & Metalloids


Molecules and Chemical Formulae

The subscript to the right of

the symbol of an element tells

the number of atoms of that

element in one molecule of

the compound.

Notice how the composition of

each compound is given by

its chemical formula.

Figure 1.29


Diatomic Molecules

These seven elements occur naturally as molecules containing two atoms.

Figure 1.28


Molecular CompoundsMolecular compounds are composed of molecules and almost

always contain only nonmetals.

Types of FormulaeEmpirical formulae give the lowest whole-number ratio of atoms

of each element in a compound, e.g. HO.

Molecular formulae give the exact number of atoms of each

element in a compound, e.g. H2O2.

Structural formulae show which atoms are attached to which

within the molecule, e.g. H-O-O-H.


Picturing Molecules

Different

representations of

the methane (CH4)

molecule.

Figure 1.30


Ions and Ionic Compounds

When atoms lose or gain electrons, they become ions.

Cations are positive and are formed by elements on the left side of the periodic chart.

Anions are negative and are formed by elements on the right side of the periodic chart.

Figure 1.31


Ionic Compounds

Ionic compounds (such as NaCl) are

generally formed between metals and

nonmetals.

Figure 1.32


Writing Formulae

Because compounds are electrically neutral, one

can determine the formula of a compound by:

writing the value of the charge on the cation

as the subscript on the anion.

writing the value of the charge on the anion

as the subscript on the cation.

Note: if the subscripts are not in the lowest

whole number ratio, simplify it, e.g. Ca2O2

would become CaO.


Chemical NomenclaturePositive Ions (Cations)

a) Cations formed from metal atoms have the same

name as the metal, e.g. Na+ is the sodium ion.

b) If a metal can form different cations, the positive

charge is indicated by a Roman numeral in

parentheses following the name of the metal,

e.g. Au+ is the gold(I) ion and Au3+ is the gold(III)

ion.

c) Cations formed from nonmetal atoms have

names that end in -ium, e.g. NH4+ is the

ammonium ion.


Chemical NomenclatureCommon Cations

Table 1.8


Chemical NomenclatureNegative Ions (Anions)

a) The names of the monatomic anions are formed

by replacing the ending of the name of the element

with -ide, e.g. O2- is the oxide ion.

b) Polyatomic anions containing oxygen (called

oxyanions) have names ending in -ate or -ite, e.g.

SO42- is the sulfate ion and SO3

2- is the sulfite ion.

c) Anions derived by adding H+ to an oxyanion are

named by adding the prefix hydrogen or

dihydrogen, e.g. HCO3-is the hydrogen carbonate

ion.


Chemical NomenclatureCommon Anions

Table 1.9


Chemical NomenclatureMore on naming oxyanions

Examples:

ClO4-

perchlorate ion (one more O atom than chlorate)

ClO3-

chlorate

ClO2-

chlorite ion (one less O atom than chlorate)

ClO-

hypochlorite ion (one O atom less than chlorite)

Names of ionic compounds consist of the cation followed by the

anion name, e.g. Cu(ClO4)2 is copper(II) perchlorate, and CaCO3

is calcium carbonate.


Chemical NomenclatureName and Formulae of Acids

1. Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid.

2. Acids containing anions whose names end in -ate or -ite are named by changing the -ate ending to -ic and the -ite ending to -ous and then adding the word acid.

Figure 1.36


Chemical NomenclatureBinary Molecular Compounds

1. The name of the element farther to the left in the periodic table is written first.

2. If both elements are in the same group in the periodic table, the one having the higher atomic number is written first.

3. The name of the second element is given an -ide ending.

4. Greek prefixes are used to indicate the number of atoms of each element.

Examples

N2O4 is dinitrogen tetroxide

P4S10 is tetraphosphorus decasulfide.

Table 1.10

Documents

Introduction: Matter, Measurement and Moleculescheminnerweb.ukzn.ac.za/Files/Chem 110 (2012)/Chem 110 2012 Ch1... · “study of matter & changes it undergoes ... Changes in matter