Lecture 1- Matter & Measurement
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General Chemistry Course # 111, two credits Second Semester 2009
King Saud bin Abdulaziz University for Health Science Textbook: Principles of Modern Chemistry by David W. Oxtoby, H. Pat Gillis, and Alan Campion (6 edition; 2007)
Dr. Rabih O. Al-Kaysi Ext: 47247 Email: [email protected]
Lecture 1 Matter & Measurement
: Very important, might be tested on
1 - Why Study Chemistry? • Chemistry is the study and investigation of the properties of matter and the chemical changes that matter undergoes. • Chemistry is one of the basic physical sciences. • Chemistry is central to our understanding of other sciences. Physical chemistry, Organic chemistry, Inorganic chemistry, Nuclear Chemistry, Analytical chemistry, Biochemistry, Nano chemistry, Astro-chemistry…. • Chemistry is all around us: Eating an apple, cooking, photosynthesis, gun powder. “Chemistry = life” • Without chemistry as a science we would not have a modern society, civilization or medicine.
The Scientific Method
Chemistry is a science ⇒ Must follow the scientific method.
What is the Scientific Method? 1- Observe a phenomena “Burning a peace of wood” 2- Experiment under controlled conditions “Burn a peace of wood in a closed container under constant temperature and pressure, the peace of wood burns then gets extinguished” 3- Theorize: Come up with a theory which explains the observed phenomena “An invisible gas in air helps sustain combustion, the same gas helps sustain life as well” 4- Test and retest the theory. Ability to predict the outcome of further experiments. “Place a mouse in the jar with the extinguished flame a see if the mouse survives”
2 - The Study of Chemistry
• Matter = physical material of the universe. • Matter is made up of few elements in an infinite combination • On the microscopic level, matter consists of atoms & molecules • Atoms combine to form molecules. Molecules combine to form DNA and life. • Molecules may consist of the same type of atoms or different types of atoms. e.g.: O2, H2O
3 - Molecular Perspective of Chemistry
4 - Classification of Matter States of Matter Three states of matter: gas, liquid, or solid. Gases: Take the shape and volume of their container. Gases: Can be compressed to form liquids or solids Liquids: Take the shape of their container, but they do have their own volume. • Solids: Are rigid and have a definite shape and volume. • • • •
Classification of Matter • Pure Substances and Mixtures • If matter is not uniform throughout (on a molecular level), then it is a heterogeneous mixture. • If matter is uniform throughout (on a molecular level), it is homogeneous. • Homogeneous matter: can be separated by physical means, then the matter is a mixture. • Homogeneous matter: cannot be separated by physical means, then the matter is a pure substance. • Pure substance: can be decomposed into something else, then the substance is a compound.
• Pure Substances and Mixtures
Classification of Matter • • • • •
Pure Substances and Mixtures If matter is not uniform throughout, then it is a heterogeneous mixture. Mixture of salt and pepper If matter is uniform throughout, it is homogeneous. Air If homogeneous matter can be separated by physical means, then the matter is a mixture. If homogeneous matter cannot be separated by physical means, then the matter is a pure substance. If a pure substance can be decomposed into something else, then the substance is a compound.
Classification of Matter Elements • If a pure substance cannot be decomposed into something else, then the substance is an element. • There are 114 elements known (91 naturally occurring) • Each element is given a unique chemical symbol (one or two letters). • Elements are the building blocks of matter. • The earth’s crust consists of 5 main elements. • The human body consists mostly of 3 main elements.
Classification of Matter Elements
Classification of Matter Elements • Chemical symbols with one letter have that letter capitalized (e.g., H = Hydrogen, B = Boron, C = Carbon, N = Nitrogen, etc.) • Chemical symbols with two letters have only the first letter capitalized (e.g., He = Helium , Be = Beryllium ). • Some names of elements are derived from Greek names. (e.g. Fe = Ferrum or Iron, Au = Aurum or Gold)
Classification of Matter Compounds • Most elements interact to form compounds. Example, H2O, C6H6 • The proportions of elements in compounds are the same irrespective of how the compound was formed. • Law of Constant Composition (or Law of Definite Proportions): – The composition of a pure compound is always the same regardless of the method it was synthesized. (e.g. H2O has the same composition if it was made by living creatures or in outer space)
Classification of Matter Compounds • If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas. H2O 2H + O • Pure substances that cannot be decomposed are elements. e.g. Heating Au to very high temperatures (Melts to a Liquid then Gas then Plasma)
Classification of Matter Mixtures • Heterogeneous mixtures are not uniform throughout. ( e.g. A peace of rock) • Homogeneous mixtures are uniform throughout. (e.g. salt water, clay: is a homogeneous mixture of water + metal oxides and silicates) • Homogeneous mixtures are called solutions.
5 - Properties of Matter Physical vs. Chemical Properties
• Physical properties: can be measure without changing the basic identity of the substance (e.g., color, density, odor, melting point) • Chemical properties: describe how substances react or change to form different substances (e.g., hydrogen burns in oxygen) • Intensive physical properties do not depend on how much of the substance is present. – Examples: density, temperature, and melting point.
• Extensive physical properties depend on the amount of substance present. – Examples: mass, volume, pressure.
Properties of Matter Physical and Chemical Changes
• When a substance undergoes a physical change, its physical appearance changes. - Ice melts to water then water boils to water vapor. • Physical changes do not result in a change of elemental composition. Ice has same % of H and O as water vapor. • Chemical change occurs when a substance changes its composition - TNT explodes to produce CO2 and H2O.
Properties of Matter Physical and Chemical Changes
Properties of Matter Separation of Mixtures • Mixtures can be separated if their physical properties are different. • Heterogeneous mixture: Solids can be separated from liquids by means of filtration. The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask. • You can also use a centrifuge to separate blood cells from plasma.
Separation of Mixtures
6 - Units of Measurement SI Units • There are two types of units: – fundamental (or base) units; – derived units.
• There are 7 base units in the SI system.
Units of Measurement Base SI Units
Units of Measurement SI Units Selected Prefixes used in SI System
Class Practice Examples • What is the name given to the unit that equals (a) 109 grams; (b) 10-6 second; (c) 10-3 meter • What fraction of a meter is a nanometer?
Units of Measurement SI Units • Note the SI unit for length is the meter (m) whereas the SI unit for mass is the kilogram (kg). – 1 kg weighs 2.2046 lb.
Temperature There are three temperature scales: • Kelvin Scale – – – –
Used in science. Same temperature increment as Celsius scale. Lowest temperature possible (absolute zero) is zero Kelvin. Absolute zero: 0 K = −273.15 oC.
Units of Measurement Temperature • Celsius Scale – Also used in science. – Water freezes at 0 oC and boils at 100 oC. – To convert: K = oC + 273.15.
• Fahrenheit Scale – Not generally used in science. – Water freezes at 32 oF and boils at 212 oF. – To convert:
5 °C = ( °F - 32 ) 9
9 °F = ( °C ) + 32 5
Class Practice Example • Make the following temperature conversions: (a) 68 oF to oC; (b) -36.7 oC to oF
5 °C = ( °F - 32 ) 9
9 °F = ( °C ) + 32 5
Units of Measurement Temperature
Units of Measurement Volume
• The units for volume are given by (units of length)3. – SI unit for volume is 1 m3.
• We usually use 1 mL = 1 cm3. • Other volume units: – 1 L = 1 dm3 = 1000 cm3 = 1000 mL.
Units of Measurement Volume
Units of Measurement Density • Used to characterize substances. • Defined as mass divided by volume: mass Density = volume • Units: g/cm3. • Originally based on mass (the density was defined as the mass of 1.00 g of pure water).
Class Practice Examples • Answer the following problems: • (a) Calculate the density of mercury if 1.0 x 102 g occupies a volume of 7.36 cm3. • (b) Using the density for mercury, calculate the mass of 65.0 cm3 of mercury.
7 - Uncertainty in Measurement • All scientific measures are subject to error. • These errors are reflected in the number of figures reported for the measurement. • These errors are also reflected in the observation that two successive measures of the same quantity are different. Precision and Accuracy • Measurements that are close to the “correct” value are accurate. • Measurements that are close to each other are precise.
Precision and Accuracy
Uncertainty in Measurement Significant Figures • The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device. • All the figures known with certainty plus one extra figure are called significant figures. • In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).
Uncertainty in Measurement Significant Figures • Non-zero numbers are always significant. (e.g. 1, 5678, 453.2342) • Zeros between non-zero numbers are always significant. (e.g.: 10023, 12.00213) • Zeros before the first non-zero digit are not significant. (e.g.: 0.0003 has one significant figure.) • Zeros at the end of the number after a decimal place are significant. (e.g.: 22.00, 1.000) • Zeros at the end of a number before a decimal place are ambiguous (e.g. 10,300 g).
Matter and Energy Energy is not created nor destroyed 1) Matter and energy are related by the famous E=MC2. This formula applies for nuclear reactions (Fission and Fusion). 2) Energy takes different forms. Heat, light (electromagnetic waves, UV, X-ray, Vis), sound 3) Energy of a chemical reaction is produced due to chemical bonds breaking or forming & crystals forming or melting. No matter is lost in the process (E=MC2) does not apply. The energy in = energy out. 4) Energy of a chemical reaction depends on the nature of the matter involved . Burning fuel produces energy because fuel is reacting with oxygen in the air to break the hydrocarbon bonds and produce water vapor, CO2 and energy in the form of light and heat. Some chemical reactions absorb energy: Decomposition of carbonic acid (H2CO3) to give CO2 and H2O. 5) Chemical reactions that liberate heat energy are called (EXOTHERMIC) Chemical reactions that absorb heat energy are called (ENDOTHERMIC) The human body functions by using exothermic chemical reaction.
King Saud bin Abdulaziz University for Health Science Textbook: Principles of Modern Chemistry by David W. Oxtoby, H. Pat Gillis, and Alan Campion (6 edition; 2007)
Dr. Rabih O. Al-Kaysi Ext: 47247 Email: [email protected]
Lecture 1 Matter & Measurement
: Very important, might be tested on
1 - Why Study Chemistry? • Chemistry is the study and investigation of the properties of matter and the chemical changes that matter undergoes. • Chemistry is one of the basic physical sciences. • Chemistry is central to our understanding of other sciences. Physical chemistry, Organic chemistry, Inorganic chemistry, Nuclear Chemistry, Analytical chemistry, Biochemistry, Nano chemistry, Astro-chemistry…. • Chemistry is all around us: Eating an apple, cooking, photosynthesis, gun powder. “Chemistry = life” • Without chemistry as a science we would not have a modern society, civilization or medicine.
The Scientific Method
Chemistry is a science ⇒ Must follow the scientific method.
What is the Scientific Method? 1- Observe a phenomena “Burning a peace of wood” 2- Experiment under controlled conditions “Burn a peace of wood in a closed container under constant temperature and pressure, the peace of wood burns then gets extinguished” 3- Theorize: Come up with a theory which explains the observed phenomena “An invisible gas in air helps sustain combustion, the same gas helps sustain life as well” 4- Test and retest the theory. Ability to predict the outcome of further experiments. “Place a mouse in the jar with the extinguished flame a see if the mouse survives”
2 - The Study of Chemistry
• Matter = physical material of the universe. • Matter is made up of few elements in an infinite combination • On the microscopic level, matter consists of atoms & molecules • Atoms combine to form molecules. Molecules combine to form DNA and life. • Molecules may consist of the same type of atoms or different types of atoms. e.g.: O2, H2O
3 - Molecular Perspective of Chemistry
4 - Classification of Matter States of Matter Three states of matter: gas, liquid, or solid. Gases: Take the shape and volume of their container. Gases: Can be compressed to form liquids or solids Liquids: Take the shape of their container, but they do have their own volume. • Solids: Are rigid and have a definite shape and volume. • • • •
Classification of Matter • Pure Substances and Mixtures • If matter is not uniform throughout (on a molecular level), then it is a heterogeneous mixture. • If matter is uniform throughout (on a molecular level), it is homogeneous. • Homogeneous matter: can be separated by physical means, then the matter is a mixture. • Homogeneous matter: cannot be separated by physical means, then the matter is a pure substance. • Pure substance: can be decomposed into something else, then the substance is a compound.
• Pure Substances and Mixtures
Classification of Matter • • • • •
Pure Substances and Mixtures If matter is not uniform throughout, then it is a heterogeneous mixture. Mixture of salt and pepper If matter is uniform throughout, it is homogeneous. Air If homogeneous matter can be separated by physical means, then the matter is a mixture. If homogeneous matter cannot be separated by physical means, then the matter is a pure substance. If a pure substance can be decomposed into something else, then the substance is a compound.
Classification of Matter Elements • If a pure substance cannot be decomposed into something else, then the substance is an element. • There are 114 elements known (91 naturally occurring) • Each element is given a unique chemical symbol (one or two letters). • Elements are the building blocks of matter. • The earth’s crust consists of 5 main elements. • The human body consists mostly of 3 main elements.
Classification of Matter Elements
Classification of Matter Elements • Chemical symbols with one letter have that letter capitalized (e.g., H = Hydrogen, B = Boron, C = Carbon, N = Nitrogen, etc.) • Chemical symbols with two letters have only the first letter capitalized (e.g., He = Helium , Be = Beryllium ). • Some names of elements are derived from Greek names. (e.g. Fe = Ferrum or Iron, Au = Aurum or Gold)
Classification of Matter Compounds • Most elements interact to form compounds. Example, H2O, C6H6 • The proportions of elements in compounds are the same irrespective of how the compound was formed. • Law of Constant Composition (or Law of Definite Proportions): – The composition of a pure compound is always the same regardless of the method it was synthesized. (e.g. H2O has the same composition if it was made by living creatures or in outer space)
Classification of Matter Compounds • If water is decomposed, then there will always be twice as much hydrogen gas formed as oxygen gas. H2O 2H + O • Pure substances that cannot be decomposed are elements. e.g. Heating Au to very high temperatures (Melts to a Liquid then Gas then Plasma)
Classification of Matter Mixtures • Heterogeneous mixtures are not uniform throughout. ( e.g. A peace of rock) • Homogeneous mixtures are uniform throughout. (e.g. salt water, clay: is a homogeneous mixture of water + metal oxides and silicates) • Homogeneous mixtures are called solutions.
5 - Properties of Matter Physical vs. Chemical Properties
• Physical properties: can be measure without changing the basic identity of the substance (e.g., color, density, odor, melting point) • Chemical properties: describe how substances react or change to form different substances (e.g., hydrogen burns in oxygen) • Intensive physical properties do not depend on how much of the substance is present. – Examples: density, temperature, and melting point.
• Extensive physical properties depend on the amount of substance present. – Examples: mass, volume, pressure.
Properties of Matter Physical and Chemical Changes
• When a substance undergoes a physical change, its physical appearance changes. - Ice melts to water then water boils to water vapor. • Physical changes do not result in a change of elemental composition. Ice has same % of H and O as water vapor. • Chemical change occurs when a substance changes its composition - TNT explodes to produce CO2 and H2O.
Properties of Matter Physical and Chemical Changes
Properties of Matter Separation of Mixtures • Mixtures can be separated if their physical properties are different. • Heterogeneous mixture: Solids can be separated from liquids by means of filtration. The solid is collected in filter paper, and the solution, called the filtrate, passes through the filter paper and is collected in a flask. • You can also use a centrifuge to separate blood cells from plasma.
Separation of Mixtures
6 - Units of Measurement SI Units • There are two types of units: – fundamental (or base) units; – derived units.
• There are 7 base units in the SI system.
Units of Measurement Base SI Units
Units of Measurement SI Units Selected Prefixes used in SI System
Class Practice Examples • What is the name given to the unit that equals (a) 109 grams; (b) 10-6 second; (c) 10-3 meter • What fraction of a meter is a nanometer?
Units of Measurement SI Units • Note the SI unit for length is the meter (m) whereas the SI unit for mass is the kilogram (kg). – 1 kg weighs 2.2046 lb.
Temperature There are three temperature scales: • Kelvin Scale – – – –
Used in science. Same temperature increment as Celsius scale. Lowest temperature possible (absolute zero) is zero Kelvin. Absolute zero: 0 K = −273.15 oC.
Units of Measurement Temperature • Celsius Scale – Also used in science. – Water freezes at 0 oC and boils at 100 oC. – To convert: K = oC + 273.15.
• Fahrenheit Scale – Not generally used in science. – Water freezes at 32 oF and boils at 212 oF. – To convert:
5 °C = ( °F - 32 ) 9
9 °F = ( °C ) + 32 5
Class Practice Example • Make the following temperature conversions: (a) 68 oF to oC; (b) -36.7 oC to oF
5 °C = ( °F - 32 ) 9
9 °F = ( °C ) + 32 5
Units of Measurement Temperature
Units of Measurement Volume
• The units for volume are given by (units of length)3. – SI unit for volume is 1 m3.
• We usually use 1 mL = 1 cm3. • Other volume units: – 1 L = 1 dm3 = 1000 cm3 = 1000 mL.
Units of Measurement Volume
Units of Measurement Density • Used to characterize substances. • Defined as mass divided by volume: mass Density = volume • Units: g/cm3. • Originally based on mass (the density was defined as the mass of 1.00 g of pure water).
Class Practice Examples • Answer the following problems: • (a) Calculate the density of mercury if 1.0 x 102 g occupies a volume of 7.36 cm3. • (b) Using the density for mercury, calculate the mass of 65.0 cm3 of mercury.
7 - Uncertainty in Measurement • All scientific measures are subject to error. • These errors are reflected in the number of figures reported for the measurement. • These errors are also reflected in the observation that two successive measures of the same quantity are different. Precision and Accuracy • Measurements that are close to the “correct” value are accurate. • Measurements that are close to each other are precise.
Precision and Accuracy
Uncertainty in Measurement Significant Figures • The number of digits reported in a measurement reflect the accuracy of the measurement and the precision of the measuring device. • All the figures known with certainty plus one extra figure are called significant figures. • In any calculation, the results are reported to the fewest significant figures (for multiplication and division) or fewest decimal places (addition and subtraction).
Uncertainty in Measurement Significant Figures • Non-zero numbers are always significant. (e.g. 1, 5678, 453.2342) • Zeros between non-zero numbers are always significant. (e.g.: 10023, 12.00213) • Zeros before the first non-zero digit are not significant. (e.g.: 0.0003 has one significant figure.) • Zeros at the end of the number after a decimal place are significant. (e.g.: 22.00, 1.000) • Zeros at the end of a number before a decimal place are ambiguous (e.g. 10,300 g).
Matter and Energy Energy is not created nor destroyed 1) Matter and energy are related by the famous E=MC2. This formula applies for nuclear reactions (Fission and Fusion). 2) Energy takes different forms. Heat, light (electromagnetic waves, UV, X-ray, Vis), sound 3) Energy of a chemical reaction is produced due to chemical bonds breaking or forming & crystals forming or melting. No matter is lost in the process (E=MC2) does not apply. The energy in = energy out. 4) Energy of a chemical reaction depends on the nature of the matter involved . Burning fuel produces energy because fuel is reacting with oxygen in the air to break the hydrocarbon bonds and produce water vapor, CO2 and energy in the form of light and heat. Some chemical reactions absorb energy: Decomposition of carbonic acid (H2CO3) to give CO2 and H2O. 5) Chemical reactions that liberate heat energy are called (EXOTHERMIC) Chemical reactions that absorb heat energy are called (ENDOTHERMIC) The human body functions by using exothermic chemical reaction.
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