Chapter 1 Lecture- Matter & Measurement

Chemistry

Chemistry

• is the study of properties of materials and changes that they undergo.

Chemistry

• is the study of properties of materials and changes that they undergo. • can be applied to all aspects of life (e.g., development of pharmaceuticals, leaf color change in fall, etc.).

The Atomic and Molecular Perspective of Chemistry

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter:

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter: • is the physical material of the universe.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter: • is the physical material of the universe.• has mass.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter: • is the physical material of the universe.• has mass.• occupies space.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter: • is the physical material of the universe.• has mass.• occupies space. • ~100 elements constitute all matter.

The Atomic and Molecular Perspective of Chemistry

Chemistry involves the study of the properties and the behavior of matter.

Matter: • is the physical material of the universe.• has mass.• occupies space. • ~100 elements constitute all matter.• A property is any characteristic that allows us to

recognize a particular type of matter and to distinguish it from other types of matter.

Elements

Elements

Elements

• are made up of unique atoms, the building

blocks of matter.

Elements

• are made up of unique atoms, the building

blocks of matter. • Names of the elements are derived from a

wide variety of sources (e.g., Latin or Greek, mythological characters, names of people or places).

Elements

• are made up of unique atoms, the building

blocks of matter. • Names of the elements are derived from a

wide variety of sources (e.g., Latin or Greek, mythological characters, names of people or places).

• Memorize element symbols

Molecules

Molecules

Molecules

• are combinations of atoms held together in

specific shapes.

Molecules

• are combinations of atoms held together in

specific shapes.• Macroscopic (observable) properties of matter

relate to submicroscopic realms of atoms.

Molecules

• are combinations of atoms held together in

specific shapes.• Macroscopic (observable) properties of matter

relate to submicroscopic realms of atoms.• Properties relate to composition (types of atoms

present) and structure (arrangement of atoms) present.

1.2 Classifications of Matter

Matter is classified by state (solid, liquid or gas) or by composition (element, compound or mixture).

States of Matter

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

Properties described on the macroscopic level:

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

Properties described on the macroscopic level:• gas (vapor): no fixed volume or shape, conforms

to shape of container, compressible.

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

Properties described on the macroscopic level:• gas (vapor): no fixed volume or shape, conforms

to shape of container, compressible.• liquid: volume independent of container, no fixed

shape, incompressible.

States of Matter

Solids, liquids and gases are the three forms of matter called the states of matter.

Properties described on the macroscopic level:• gas (vapor): no fixed volume or shape, conforms

to shape of container, compressible.• liquid: volume independent of container, no fixed

shape, incompressible.• solid: volume and shape independent of

container, rigid, incompressible.

States of Matter

States of Matter

Properties described on the molecular level:

States of Matter

Properties described on the molecular level:• gas: molecules far apart, move at high

speeds, collide often.

States of Matter

Properties described on the molecular level:• gas: molecules far apart, move at high

speeds, collide often.• liquid: molecules closer than gas, move

rapidly but can slide over each other.

States of Matter

Properties described on the molecular level:• gas: molecules far apart, move at high

speeds, collide often.• liquid: molecules closer than gas, move

rapidly but can slide over each other.• solid: molecules packed closely in definite

arrangements.

Pure Substances

Pure Substances

Pure substances:

Pure Substances

Pure substances: • are matter with fixed compositions and

distinct proportions.

Pure Substances

Pure substances: • are matter with fixed compositions and

distinct proportions.• are elements (cannot be decomposed into

simpler substances, i.e. only one kind of atom) or compounds (consist of two or more elements).

Mixtures

Mixtures

• are a combination of two or more pure substances.

Mixtures

• are a combination of two or more pure substances.

• Each substance retains its own identity.

Elements

Elements

• There are 116 known elements.

Elements

• There are 116 known elements.• They vary in abundance.

Elements

• There are 116 known elements.• They vary in abundance.• Each is given a unique name and is

abbreviated by a chemical symbol.

Elements

• There are 116 known elements.• They vary in abundance.• Each is given a unique name and is

abbreviated by a chemical symbol.• they are organized in the periodic table.

Elements

• There are 116 known elements.• They vary in abundance.• Each is given a unique name and is

abbreviated by a chemical symbol.• they are organized in the periodic table.• Each is given a one- or two-letter symbol

derived from its name.

Compounds

Compounds

• Compounds are combinations of elements.

Compounds

• Compounds are combinations of elements. Example: The compound H2O is a

combination of the elements H and O.

Compounds

• Compounds are combinations of elements. Example: The compound H2O is a

combination of the elements H and O.• The opposite of compound formation is

decomposition.

Compounds

• Compounds are combinations of elements. Example: The compound H2O is a

combination of the elements H and O.• The opposite of compound formation is

decomposition.• Compounds have different properties than their

component elements (e.g., water is liquid, but hydrogen and oxygen are both gases at the same temperature and pressure).

Law of Constant (Definite) Proportions

Law of Constant (Definite) Proportions

(Proust): A compound always consists of the same combination of elements (e.g., water is always 11% H and 89% O).

Mixtures

Mixtures

• A mixture is a combination of two or more pure substances.

Mixtures

• A mixture is a combination of two or more pure substances.

• Each substance retains its own identity; each substance is a component of the mixture.

Mixtures

• A mixture is a combination of two or more pure substances.

• Each substance retains its own identity; each substance is a component of the mixture.

• Mixtures have variable composition.

Mixtures

• A mixture is a combination of two or more pure substances.

• Each substance retains its own identity; each substance is a component of the mixture.

• Mixtures have variable composition.• Heterogeneous mixtures do not have uniform

composition, properties, and appearance, e.g., sand.

Mixtures

• A mixture is a combination of two or more pure substances.

• Each substance retains its own identity; each substance is a component of the mixture.

• Mixtures have variable composition.• Heterogeneous mixtures do not have uniform

composition, properties, and appearance, e.g., sand.

• Homogeneous mixtures are uniform throughout, e.g., air; they are solutions.

1.3 Properties of Matter

1.3 Properties of MatterEach substance has a unique set of physical and

chemical properties.

1.3 Properties of MatterEach substance has a unique set of physical and

chemical properties.• Physical properties are measured without

changing the substance (e.g., color, density, odor, melting point, etc.).

1.3 Properties of MatterEach substance has a unique set of physical and

chemical properties.• Physical properties are measured without

changing the substance (e.g., color, density, odor, melting point, etc.).

• Chemical properties describe how substances react or change to form different substances (e.g., hydrogen burns in oxygen).

1.3 Properties of Matter

1.3 Properties of Matter

Properties may be categorized as intensive or extensive.

1.3 Properties of Matter

Properties may be categorized as intensive or extensive.

• Intensive properties do not depend on the amount of substance present (e.g., temperature, melting point etc.).

1.3 Properties of Matter

Properties may be categorized as intensive or extensive.

• Intensive properties do not depend on the amount of substance present (e.g., temperature, melting point etc.).

• Extensive properties depend on the quantity of substance present (e.g., mass, volume etc.).

1.3 Properties of Matter

Properties may be categorized as intensive or extensive.

• Intensive properties do not depend on the amount of substance present (e.g., temperature, melting point etc.).

• Extensive properties depend on the quantity of substance present (e.g., mass, volume etc.).

• Intensive properties give an idea of the composition of a substance whereas extensive properties give an indication of the quantity of substance present.

Physical and Chemical Changes

Physical and Chemical Changes

• Physical change: substance changes physical appearance without altering its identity (e.g., changes of state).

Physical and Chemical Changes

• Physical change: substance changes physical appearance without altering its identity (e.g., changes of state).

• Chemical change (or chemical reaction): substance transforms into a chemically different substance (i.e. identity changes, e.g., decomposition of water, explosion of nitrogen triiodide).

Separation of Mixtures

Separation of Mixtures

Key: separation techniques exploit differences in properties of the components.

Separation of Mixtures

Key: separation techniques exploit differences in properties of the components.

• Filtration: remove solid from liquid.

Separation of Mixtures

Key: separation techniques exploit differences in properties of the components.

• Filtration: remove solid from liquid. • Distillation: boil off one or more

components of the mixture.

Separation of Mixtures

Key: separation techniques exploit differences in properties of the components.

• Filtration: remove solid from liquid. • Distillation: boil off one or more

components of the mixture. • Chromatography: exploit solubility of

components.

The Scientific Method

The Scientific MethodThe scientific method provides guidelines for the

practice of science.

The Scientific MethodThe scientific method provides guidelines for the

practice of science. • Collect data (observe, experiment, etc.).

The Scientific MethodThe scientific method provides guidelines for the

practice of science. • Collect data (observe, experiment, etc.). • Look for patterns, try to explain them, and

develop a hypothesis or tentative explanation.

The Scientific MethodThe scientific method provides guidelines for the

practice of science. • Collect data (observe, experiment, etc.). • Look for patterns, try to explain them, and

develop a hypothesis or tentative explanation. • Test hypothesis, then refine it.

The Scientific MethodThe scientific method provides guidelines for the

practice of science. • Collect data (observe, experiment, etc.). • Look for patterns, try to explain them, and

develop a hypothesis or tentative explanation. • Test hypothesis, then refine it. • Bring all information together into a scientific

law (concise statement or equation that summarizes tested hypotheses).

The Scientific MethodThe scientific method provides guidelines for the

practice of science. • Collect data (observe, experiment, etc.). • Look for patterns, try to explain them, and

develop a hypothesis or tentative explanation. • Test hypothesis, then refine it. • Bring all information together into a scientific

law (concise statement or equation that summarizes tested hypotheses).

• Bring hypotheses and laws together into a theory. A theory should explain general principles.

Steps in the Scientific Method

1.3

Steps in the Scientific Method

1.3

Steps in the Scientific Method

1.3

Steps in the Scientific Method

1.3

Steps in the Scientific Method

1.3

Steps in the Scientific Method

1.3

Units of Measurement

SI Units

SI Units

• Système International d’Unités

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.

PRACTICE EXERCISE(a) What decimal fraction of a second is

a picosecond, ps? (b) Express the measurement 6.0 ×

103 m using a prefix to replace the power of ten.

(c) Use exponential notation to express 3.76 mg in grams.

PRACTICE EXERCISE(a) What decimal fraction of a second is

a picosecond, ps? (b) Express the measurement 6.0 ×

103 m using a prefix to replace the power of ten.

(c) Use exponential notation to express 3.76 mg in grams.

Answers: (a) 10–12 second, (b) 6.0 km, (c) 3.76 × 10–3 g

Volume

Volume

• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).

Volume

• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).□ A liter is a cube 1 dm

long on each side.

Volume

• The most commonly used metric units for volume are the liter (L) and the milliliter (mL).□ A liter is a cube 1 dm

long on each side.□ A milliliter is a cube 1 cm

long on each side.

Temperature:

A measure of the average kinetic energy of the particles in a sample.

Temperature

Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.

Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.□ 0°C is the freezing point

of water.

Temperature• In scientific

measurements, the Celsius and Kelvin scales are most often used.

• The Celsius scale is based on the properties of water.□ 0°C is the freezing point

of water.□ 100°C is the boiling

point of water.

Temperature

Temperature

• The Kelvin is the SI unit of temperature.

Temperature

• The Kelvin is the SI unit of temperature.

• It is based on the properties of gases.

Temperature

• The Kelvin is the SI unit of temperature.

• It is based on the properties of gases.

• There are no negative Kelvin temperatures.

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

Temperature

Temperature

• The Fahrenheit scale is not used in scientific measurements.

Temperature

• The Fahrenheit scale is not used in scientific measurements.

• °F = 9/5(°C) + 32

Temperature

• The Fahrenheit scale is not used in scientific measurements.

• °F = 9/5(°C) + 32• °C = 5/9(°F − 32)

Density:

Physical property of a substance

Density:

Physical property of a substance

d= mV

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(a) Calculate the density of mercury if 1.00 × 102 g occupies a volume of 7.36 cm3.

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(a) Calculate the density of mercury if 1.00 × 102 g occupies a volume of 7.36 cm3.

Solution (a) We are given mass and volume, so Equation 1.3 yields

(b) Calculate the volume of 65.0 g of the liquid methanol (wood alcohol) if its density is 0.791 g/mL.

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(b) Calculate the volume of 65.0 g of the liquid methanol (wood alcohol) if its density is 0.791 g/mL.

Solution (b) Solving Equation 1.3 for volume and then using the given mass and density gives

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(c) What is the mass in grams of a cube of gold (density = 19.32 g/ cm3) if the length of the cube is 2.00 cm?

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(c) What is the mass in grams of a cube of gold (density = 19.32 g/ cm3) if the length of the cube is 2.00 cm?

Solution

(c) We can calculate the mass from the volume of the cube and its density. The volume of a cube is given by its length cubed:

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

(c) What is the mass in grams of a cube of gold (density = 19.32 g/ cm3) if the length of the cube is 2.00 cm?

Solution

(c) We can calculate the mass from the volume of the cube and its density. The volume of a cube is given by its length cubed:

Solving Equation 1.3 for mass and substituting the volume and density of the cube, we have

SAMPLE EXERCISE 1.4 Determining Density and Using Density to Determine Volume or Mass

Uncertainty in Measurements

Different measuring devices have different uses and different degrees of accuracy.

PRACTICE EXERCISEA balance has a precision of ± 0.001 g. A sample that has a mass of about 25 g is placed on this balance. How many significant figures should be reported for this measurement?

PRACTICE EXERCISEA balance has a precision of ± 0.001 g. A sample that has a mass of about 25 g is placed on this balance. How many significant figures should be reported for this measurement?

Answer: five, as in the measurement 24.995 g

Uncertainty in Measurement

Significant Figures

Significant Figures

• The term significant figures refers to digits that were measured.

Significant Figures

• The term significant figures refers to digits that were measured.

• When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

Significant Figures

Significant Figures

1. All nonzero digits are significant.

Significant Figures

1. All nonzero digits are significant.2. Zeroes between two significant figures

are themselves significant.

Significant Figures

1. All nonzero digits are significant.2. Zeroes between two significant figures

are themselves significant.3. Zeroes at the beginning of a number

are never significant.

Significant Figures

1. All nonzero digits are significant.2. Zeroes between two significant figures

are themselves significant.3. Zeroes at the beginning of a number

are never significant.4. Zeroes at the end of a number are

significant if a decimal point is written in the number.

SAMPLE EXERCISE 1.6 Determining the Number of Significant Figures in a Measurement

How many significant figures are in each of the following numbers (assume that each number is a measured quantity): (a) 4.003, (b) 6.023 × 1023, (c) 5000?

SAMPLE EXERCISE 1.6 Determining the Number of Significant Figures in a Measurement

How many significant figures are in each of the following numbers (assume that each number is a measured quantity): (a) 4.003, (b) 6.023 × 1023, (c) 5000?

Solution (a) Four; the zeros are significant figures. (b) Four; the exponential term does not add to the number of significant figures.

SAMPLE EXERCISE 1.6 Determining the Number of Significant Figures in a Measurement

How many significant figures are in each of the following numbers (assume that each number is a measured quantity): (a) 4.003, (b) 6.023 × 1023, (c) 5000?

Solution (a) Four; the zeros are significant figures. (b) Four; the exponential term does not add to the number of significant figures. (c) One. We assume that the zeros are not significant when there is no decimal point shown. If the number has more significant figures, a decimal point should be employed or the number written in exponential notation. Thus, 5000. has four significant figures, whereas 5.00 × 103 has three.

Significant Figures

Significant Figures

• When addition or subtraction is performed, answers are rounded to the least significant decimal place.

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.

SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a Calculated Quantity

The width, length, and height of a small box are 15.5 cm, 27.3 cm, and 5.4 cm, respectively. Calculate the volume of the box, using the correct number of significant figures in your answer.

SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a Calculated Quantity

The width, length, and height of a small box are 15.5 cm, 27.3 cm, and 5.4 cm, respectively. Calculate the volume of the box, using the correct number of significant figures in your answer.

Solution In multiplication & division, count sig figs.In addition & subtraction, count decimal places.

SAMPLE EXERCISE 1.7 Determining the Number of Significant Figures in a Calculated Quantity

The width, length, and height of a small box are 15.5 cm, 27.3 cm, and 5.4 cm, respectively. Calculate the volume of the box, using the correct number of significant figures in your answer.

Solution In multiplication & division, count sig figs.In addition & subtraction, count decimal places.

A gas at 25°C fills a container whose volume is 1.05 × 103

cm3. The container plus gas have a mass of 837.6 g. The container, when emptied of all gas, has a mass of 836.2 g. What is the density of the gas at 25°C?

SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a Calculated Quantity

A gas at 25°C fills a container whose volume is 1.05 × 103

cm3. The container plus gas have a mass of 837.6 g. The container, when emptied of all gas, has a mass of 836.2 g. What is the density of the gas at 25°C?

Solution

To calculate the density, we must know both the mass and the volume of the gas. mass of the gas = full - empty container: (837.6 – 836.2) g = 1.4 g

SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a Calculated Quantity

A gas at 25°C fills a container whose volume is 1.05 × 103

cm3. The container plus gas have a mass of 837.6 g. The container, when emptied of all gas, has a mass of 836.2 g. What is the density of the gas at 25°C?

Solution

To calculate the density, we must know both the mass and the volume of the gas. mass of the gas = full - empty container: (837.6 – 836.2) g = 1.4 g

SAMPLE EXERCISE 1.8 Determining the Number of Significant Figures in a Calculated Quantity

Accuracy versus Precision

Accuracy versus Precision

• Accuracy refers to the proximity of a measurement to the true value of a quantity.

Accuracy versus Precision

• 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.

SAMPLE EXERCISE 1.9 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)

SAMPLE EXERCISE 1.9 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)

SAMPLE EXERCISE 1.9 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)

Solution Because we want to change from lb to g, we look for a relationship between these units of mass. From the back inside cover we have 1 lb = 453.6 g. In order to cancel pounds and leave grams, we write the conversion factor with grams in the numerator and pounds in the denominator:

SAMPLE EXERCISE 1.9 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)

Solution Because we want to change from lb to g, we look for a relationship between these units of mass. From the back inside cover we have 1 lb = 453.6 g. In order to cancel pounds and leave grams, we write the conversion factor with grams in the numerator and pounds in the denominator:

SAMPLE EXERCISE 1.9 Converting Units

If a woman has a mass of 115 lb, what is her mass in grams? (Use the relationships between units given on the back inside cover of the text.)

Solution Because we want to change from lb to g, we look for a relationship between these units of mass. From the back inside cover we have 1 lb = 453.6 g. In order to cancel pounds and leave grams, we write the conversion factor with grams in the numerator and pounds in the denominator:

The answer can be given to only three significant figures, the number of significant figures in 115 lb.

SAMPLE EXERCISE 1.10 Converting Units Using Two or More Conversion Factors

The average speed of a nitrogen molecule in air at 25°C is 515 m/s. Convert this speed to miles per hour.

SAMPLE EXERCISE 1.10 Converting Units Using Two or More Conversion Factors

The average speed of a nitrogen molecule in air at 25°C is 515 m/s. Convert this speed to miles per hour.

Solution On the back inside cover of the book, we find that 1 mi = 1.6093 km1 km = 103 m60 s = 1 min 60 min = 1 hr

SAMPLE EXERCISE 1.10 Converting Units Using Two or More Conversion Factors

The average speed of a nitrogen molecule in air at 25°C is 515 m/s. Convert this speed to miles per hour.

Solution On the back inside cover of the book, we find that 1 mi = 1.6093 km1 km = 103 m60 s = 1 min 60 min = 1 hr

SAMPLE EXERCISE 1.11 Converting Volume Units

Earth’s oceans contain approximately 1.36 × 109 km3 of water. Calculate the volume in liters.

SAMPLE EXERCISE 1.11 Converting Volume Units

Earth’s oceans contain approximately 1.36 × 109 km3 of water. Calculate the volume in liters.

Solution1 L = 10–3 m3

1 km = 103 m

SAMPLE EXERCISE 1.11 Converting Volume Units

Earth’s oceans contain approximately 1.36 × 109 km3 of water. Calculate the volume in liters.

Solution1 L = 10–3 m3

1 km = 103 m

Thus, converting from km3 to m3 to L, we have