1. What is CHEMISTRY ? The Study of Matter and its Properties, the Changes that Matter Undergoes,...

What is CHEMISTRY ?

The Study ofMatter and its Properties,the Changes thatMatter Undergoes, and the EnergyAssociated withthose Changes

CHAPTER

1 What is CHEMISTRY ?

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Analytical (qualitative and quantitative) Chemistry:

Qualitative Analytical Chemistry: What is a sample of matter composed of?

Quantitative Analytical Chemistry: How much “stuff” is in a sample of matter?

Physical Chemistry:

The study of the physics involved with chemical changes.

Organic Chemistry:

The study of properties and reactions of compounds that contain Carbon

Inorganic Chemistry:

The study of properties and reactions of compounds that are not Carbon based

Biochemistry: The study of living systems (Biology + Chemistry)

CHAPTER

1 Major Chemistry subdivisions

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From experiment to modelling

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“Chemists communicate their observations and ideas with each other through words, labels, drawings and symbols to refer to particular events and substances”

Levels of OperationCHAPTER

1Chemistry: The Study of Change

Classifying Matter

• Matter is defined as: Anything that occupies space and has mass

• The classification of matter include substances, mixtures, elements, and compounds, as well as atoms and molecules.

• Mass: The quantity of matter in a material.• Law of Conservation of Mass:

The total mass remains constant during a chemical change (chemical reaction).

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CHAPTER


Solid: characterized by rigidity; fixed volume and fixed shape.

Liquid: relatively incompressible fluid; fixed volume, no fixed shape.

Gas: compressible fluid; no fixed volume, no fixed shape.

States of MatterCHAPTER

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Particles are closely packedand vibrate about fixed positions

Particles are not in fixed positions andcan move past one another

Particles are not in fixed positions andcan move past one another

Chemistry: The Study of Change

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• Substance: is a single, pure form of matter– Example: water, copper, sugar, Oxygen and gold

• Or, A kind of matter that cannot be separated into other kinds of matter by any physical process such as distillation or sublimation.

• Mixtures contain two of more substances mixed together– Example: air, steel, soda-cola, saltwater, salad dressing– Mixtures can be:

• Homogeneous mixtures• Heterogeneous mixtures

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Classifying MatterCHAPTER


• Homogeneous and heterogeneous mixtures

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Noodle soup is a heterogeneous mixtureMaterials are not uniformly dispersed over the whole sample

A solution of salt in water is a homogeneous mixture or sugar dissolves in water , The sample has a uniform composition throughout



• There are only 118 (so far) known elements• 90 elements are found in nature, the rest have been

made in the laboratoryExample of elements:

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CHAPTER

1 Elements and their Atoms

Chemistry: The Study of Change

Element: A substance that cannot be decomposed into simpler substances by any chemical reaction.

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Compounds and MoleculesCHAPTER


Molecules

Compounds

NaCl

O2

H2

CaCl2

CH4

A molecule is formed when two or more atoms join together chemically.

A compound is a molecule that contains at least two different elements.

All compounds are molecules but not all molecules are compounds.

http://education.jlab.org/qa/atom.html

http://education.jlab.org/qa/element.html

• A CHEMICAL COMPOUND is a substance composed of atoms of two or more different elements. H2O or CO2

– A combination of atoms represents a MOLECULE or ION, which is characteristic for each compound

• Molecules or ions are composed of atoms in fixed and definite proportions

• Compounds have properties which are different than properties of elements they are made from

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CompoundsCHAPTER


EXAMPLE: Sodium (metal) and chlorine (gas) will react to produce sodium chloride NaCl (table salt – colourless, crystalline substance):

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CompoundsCHAPTER


• A CHEMICAL FORMULA shows the composition of a compound

• It features the relative numbers of different types of atomsEXAMPLE:

NaClSodium chloride containsthe same number of atomsof sodium and chlorinecombined in a molecule

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Na

Cl

CompoundsCHAPTER


• Kinetic-molecular model of matter– Particles: atoms, molecules or ions– Motion of particles opposed by the forces of

attraction between particles

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SOLID LIQUID GAS

Forces of attraction dominate

Fast motion of particles due to the increase in kinetic energy. Resulting forces overcome forces of attraction

Temperature increases



Aluminum powder burns in oxygen to produce a substance called aluminum oxide. A sample of 2.00 grams of aluminum is burned in oxygen and produces 3.78 grams of aluminum oxide. How many grams of oxygen were used in this reaction?

aluminum + oxygen = aluminum oxide 2.00 g + oxygen = 3.78 g

oxygen = 1.78 g

Chemical Reactions, Chemical ChangeCHAPTER


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• Classification of Matter

Matter is anything that occupies space and has mass. The classification of matter include substances, mixtures, elements, and compounds, as well as atoms and molecules.

• Substances and Mixtures:• A Substance is a form of matter that has a definite (constant)

composition and distinct properties. Examples are water, ammonia, gold and oxygen.

• A Mixture is a combination of two or more substances in the substances retain their distinct identities. Examples are air, and cement.

• Mixtures are either homogeneous or heterogeneous. When a spoonful of sugar dissolves in water we obtain a homogenous mixture in which the composition of the mixture is the same throughout. If sand is mixed with iron filings, however, the sand grains remain separate. This type of mixture is called a heterogeneous mixture because the composition is not uniform.

IntroductionCHAPTER


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Any type of mixture, can be created and then separated by physical means into pure component.

Solution: Is a homogeneous mixture and its components called :Solute : it is the existent substance in a low quantity.Solvent: it is the existent substance in a high quantity.

Elements and CompoundsAn element is a substance that cannot be separated into simpler substances. To date, 117 elements have been identified.Atoms of most elements can interact with one another to form compounds.Hydrogen gas, for example, burns in oxygen gas to form water, which has properties that are different from those of the starting materials.

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A measured quantity is usually written as a number with an appropriate unit.Examples20 grams6.63 × 10–34 joule·seconds

SI UnitsIn 1960 the General Conference of Weights and Measures, the international authority of units, proposed a revised metric system called international system of units (abbreviated SI . Table 1.2 shows the seven SI base units. All other units of measurements can be derived from these base units. SI units are modified in decimal fashion by a series of prefixes, as shown in table 1.3

Measurements CHAPTER


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femto- f 0.000000000000001 10-15 1 femtometer fm = 1 x 10-15 m22

Note that a metric prefix simply represents a number: 1mm = 1 x 10-3 mMeasurements that we will utilize frequently in our study of chemistry include time, mass, volume, and temperature.



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Mass and WeightMass is a measure of the amount of matter in an object, weight, is the force that gravity exerts on an object. An apple that falls from a tree is pulled downward by Earth`s gravity. The mass of the apple is constant and does not depend on its location, but its weight does. Chemists are interested in mass, which can be determined with a balance.

The SI unit of mass is kilogram (Kg) 1 Kg + 1000 g = 1 x 103 g

Measurement CHAPTER


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Volume V = length x width x height = m x m x m = m3

The SI unite of length is the meter (m), and that of volume is the cubic meter (m3), the cubic centimeter (cm3) and the cubic decimeter (dm3)

1cm3 = ( 1 x 10-2 m )3 = 1 x 10-6 m3

1dm3 = ( 1 x 10-1 m)3 = 1 x 10-3 m3

Another common unite of volume is the liter (L)1L = 1000 ml = 1000 cm3

= 1 dm3

1 mL = 1 cm3.

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DensityThe equation for density is

The SI-derived unit for density is the kilogram per cubic meter ( kg/m3 ). This unit is large for most chemical applications. Therefore, gram per cubic centimeter (g/cm3) and its equivalent, gram per milliliter (g/mL) are used for solid and liquid densities. 1g/cm3 = 1g/mL = 1000kg/m3 Example 1.1A piece of metal with a mass of 200 g has a volume of 17.6 cm3 . Calculate the density of gold. Solution:We are given the mass and volume and asked to calculate the density

Example 1.2The density of mercury, the only metal is a liquid at room temperature, is 13.6 g/ml.calculate the mass of 7.8 mL of the liquid.Solution:

density = mass

volume, d =

m

V

d = m

V

13.6 X 7.8 = 106.08 gm = d X V =

d=20017.6

= 11.36 g/cm3

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Heat and Temperature

• Heat and Temperature are not the same thingT is a measure of the intensity of heat in a body

• 3 common temperature scales - all use water as a reference

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Temperature:Three temperature are currently in use. Their units are ⁰F(degrees Fahrenheit). ⁰C (degree s Celsius), and K (kelvin). The fahrenheit scale, defines the normal freezing and boiling points of water to be 32⁰F and 212⁰F, respectively. The Celsius scale divides the range between the freezing point (0⁰C) and boiling point (100⁰C) of water into 100 degrees. The kelvin is the SI base unite of temperature: it is the absolute temperature scale. T e zero on kelvin ( 0 K ) is the lowest temperature can be attained. To convert Fahrenheit we use

K = ⁰C + 273,15

C = ( F - 32 F) x 5 C9 F

oo o oo

?

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Heat and Temperature

MP water BP water• Fahrenheit 32 oF 212 oF• Celsius 0.0 oC 100 oC• Kelvin 273 K 373 K

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Relationships of the Three Temperature Scales

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Easy method to remember how to convert from Centigrade to Fahrenheit.

1. Double the Centigrade temperature.2. Subtract 10% of the doubled number.3. Add 32.

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Temperature

Example: Express 548 K in Celsius degrees.

Example: Convert 211oF to degrees Celsius.

273KCo

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Example 1.3a) A certain solder (alloy of Tin & lead) has a melting point of 245⁰C. What is its melting point in degrees Fahrenheit?b) Helium has the lowest boiling points of all the elements at -396⁰F. Convert this temperature to degrees Celsius.c) Mercury, the only metal that exists as a liquid at room temperature, melts at -96.5⁰C , convert its melting point to kelvins.Solution:a) 245⁰C = ( ⁰F – 32⁰F) x

(245⁰C x 9⁰F ÷ 5⁰C ) + 32⁰F = ⁰F

(245x 9⁰F ÷ 5 ) + 32⁰ F =473 ⁰F

c) K = C + 273,15 = -38.9 + 273.15 = 234.3 K⁰

Other examples in “ summery of G. Chemistry” page 10

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Handling NumbersScientific NotationAll numbers can be expressed in the form N x 10n

where N is a number between 1 and 10 and n, the exponent, is a positive or negative integer(whole number). Any number expressed in this way is said to be written in scientific notation.e.g. 1) Express 568.762 in scientific notation 568.762 = 5.68762 x 102

2) Express 0.00000772 in scientific notation 0.00000772 = 7.72 x 10-6

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Significant Figures• Any digit that is not zero is significant

– 1.234 kg 4 significant figures• Zeros between nonzero digits are significant

– 606 m 3 significant figures– 16.07 4 sig figs.

• Zeros to the left of the first nonzero digit are not significant– 0.08 L 1 significant figure

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Significant Figures• If a number is greater than 1, then all zeros to

the right of the decimal point are significant– 2.0 mg 2 significant figures

• If a number is less than 1, then only the zeros that are at the end and in the middle of the number are significant– 0.00420 g 3 significant figures

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Trailing zeros are zeros at the right end of the number. They are significant only if the number contains a decimal point. 9.300 has 4 sig figs. 150 has 2 sig figs. 1000000 1 sig figs. 1.000000 7 sig fig

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Significant Figures

• Exact numbers: have an INFINITE number of significant figures. This rule applies to numbers that are definitions.

• For example, 1 meter = 1.00 meters = 1.0000 meters = 1.0000000000000000000 meters, etc.

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