CHEM 1070

Prelude - Survey of Chemistry

Chemistry

  • study of matter and its properties

  • the science of molecules and their transformations

        centrally important to all of science such as medicine, agriculture, industry,

        and materials

    Humans are part of chemistry - bodies-organs-tissues-cells (site of elaborate

    chemical processes)

Modern Chemistry

  • Organic Chemistry - study of substances containing carbon

  • Inorganic Chemistry - study of substances which do not contain carbon

  • Biochemistry - study of substances found in plants and animals

  • Analytical Chemistry - the branch of chemistry dealing with the quantitative and qualitative analysis of substances

  • Physical Chemistry - the branch of chemistry dealing with the mathematical and physical behavior of substances

Brief History of Chemistry

  • The Stone Age (8000 BC)

    • Early men used materials as they found them (rocks, wood, bone)

  • Metals (4000 BC)

    • Men found enough free metals and were impressed enough with their properties to seek more of them

           malleable - tools could be made and sharpened (copper and gold)

    • Around 3000 BC nuggets were no longer the only source of metals

                    blue rocks (ores) + fire → copper

                    copper + tin → bronze

  • The Bronze Age (3000BC - 1100BC)

    • Also known for a better metal - Iron

    • Couldn’t be extracted from ores and couldn’t get fires hot enough

    • 1500BC Empire of Asia Minor (modern day Turkey)

         Iron smelting - required bellows and charcoal

  • The Iron Age (900BC)

    • The emergence of the practical chemical arts

    • Egyptians used pigments and extracts from plants for artistic and religious purposes

  • Greek ‘Elements’ (600BC) - The first chemical theorists

    • What was everything made of? (Thales 640 - 546BC)

      • everything is water

    • Anaximenes (585 - 528BC)

      • air is the element of the universe

    • Heraclitus (540 - 475BC)

      • Said change characterizes the universe (element is fire)

    • Empedocles (490 - 430BC)

      • Compromised: could there be more than one element

            water, air, fire, earth

      • Doctrine of the four elements

  • Aristotle (384 - 322BC)

    • Accepted the notion of the four elements and its holds swayed over men’s thinking for over 2,000 years

  • Democritus (470 - 380BC)

    • Believed the matter is not infinitely divisible, called the smallest particle “atomos” meaning indivisible

  • Accelerators - Atom Smasher

    • LHS

    • Fermilab

  • Age of Alchemy (300BC - 1600AD)

    • Sought gold from base metals - transmutations

  • Some progress in the 17th Century

    • Comes from an unlikely source - the study of gases

  • Jan Baptista Van Helmont (Belgium 1577 - 1644)

    • Studied the vapors of burned wood

          used the Greek word ‘chaos’ to describe these vapors

  • Robert Boyle (Irish 1627 - 1691)

    • Measured the relationship between the pressure exerted by a gas and the volume it occupies

    • Boyle’s Law - pressure and volume of a gas are inversely proportional

         as the pressure increases the volume decreases

  • Joseph Priestly (Eng. 1733 - 1804)

    • Heats mercury and lead in air to obtain red powders, discovers oxygen

  • Daniel Rutherford (Eng. 1749 - 1819

    • Discovers nitrogen

  • Joseph Black (1729 - 1799)

    • Air is not an element, it contains oxygen, nitrogen, and carbon dioxide

  • Henry Cavendish (1731 - 1810)

    • Showed that water is the product reaction of specific amount of hydrogen with a specific amount of oxygen

  • Antoine Lavoisier (1743 - 1794)

    • Father of Modern Chemistry

    • Measures all things chemical

    • Theories of gas

          first to systematize naming

          no more doctrine of the four elements

Chapter I - Matter and Measurement

Measurements

    Physical quantities                Name                Abbreviation

    Mass                                    gram                     g

    Length                                  meter                   m

    Volume                                 liter                       l

    Time                                     second                 s

Metric System

  • All units are expressed as powers of 10

Selected prefixes in the metric system

  • Kilo (k)                    103                    1kg - 1000kg

  • Centi (C) 10-2 100cg - 1g

  • Milli (m) 10-3 1000mg - 1g

  • Micro (u) 10-6 1.0 × 106 ug - 1g

Problem Solving

    Dimensional analysis - a way of solving conversion problems

  • Every number in a problem must have units associated with it

  • When correctly manipulated, starting units cancel and leave the final answers in the appropriate units

  • Use conversion factor - multiplies that relate the desired unit to the starting unit

            Conversion factor - (desired unit / starting unit)

            The numerator of the conversion factor must be equivalent to the denominator

Significant Figures

    Reporting figures accurately given the measuring devices

    Rules:

  • Read the numbers from left to right starting with the first digit that is not zero

        ex.     Number                    # of sig figs

                    1.23                        3

                    0.123                      3

                    0.0023                    3

    if the number is > 1, all zeros to the right of the decimal point are significant

    if the number is < 1, all zeros to the right of the first sig fig are significant

        ex.    Number                    # of sig figs

                    2.000                     4

                    0.020                      2

    ‘trailing zeros’ may or may not be significant

        ex.    Number                    # of sig figs

                    100                         1

    a decimal point makes this unambiguous

                    100.                        3

                    100.00                    5

    (all none-zeros are significant. zeros between sig figs are significant)

  • When adding or subtracting the numbers of decimal places in the answer should be qual to the number of decimal places in the quantity with the fewest places

          ex.       0.12                      2 s.f.

                        1.6                        2 s.f.

               +     10.976                    5 s.f.

                      12.696 → answer is 12.7 since 1.6 has fewest places

  • In multiplication and division, the number of sig figs in the answer should be the same as the quantity with the fewer sig figs

            if the numerator has 4 s.f. the denominator has 3 s.f. the answer should have 3 s.f.

  • When a number is rounded off, the last digit retained is increased by 1 only if the following digit is 5 or greater

  • Exact numbers (e.g. 100 dollar bills) and defined quantities (e.g. 1000g → 1kg) have infinite number of significant figures

            (1kg / 1000g) has an infinite number of sig figs)

Prefixes used in the Metric System

    Prefix                    Abbreviation                  Value

     Mega                    M                                  106

     Kilo                       k                                   103

      Deci                      d                                   10-1

     Centi                     c                                   10-2

     Milli                       m                                  10-3

      Micro                    μ (‘mew’)                       10-6

      Nano                    n                                    10-9

     Pico                      p (‘peako’)                     10-12

Length

  • We use the meter for unit of length

        1 meter → 3.3 feet or 39.4 inches

            “a human leg is about a meter long”

        Centimeters (cm: 10-2) and meters are convenient for objects in the lab

        Nanometer (nm: 10-9) and picometers (pm: 10-12) are used for dimension at the         molecular level

Area and Volume

  • The units for area and volume are derived from the base unit of length. Area can be given in square centimeters (cm2) and volume in cubic centimeters (cm2)

        “Cubic centimeter or c.c. and mL are used interchangeably”

Volume by calculation

  • Volume of rectangular solid → l x w x h

        To calculate volume, the units must all be the same

        In the metric system, the basic unit of volume is liter

                1 liter → 10cm x 10cm x 10cm

                1 liter → 1000cm3 → 1000mL

        Volume of liquid are simply measured using calibrated glassware: breakers,     pipettes, burets, etc.

                cube → s3

                sphere →43πr3\frac{4}{3\pi r^3}

                cylinder → πr22h\pi r^22h

        The difference between the final volume and the initial volume represents the     volume of the object

                displacement → volumefinal - volumeinitial

                                    ex. 5.00 mL - 3.40 mL → 1.60 mL

Mass and Weight

  • “Mass and weight used interchangeably in everyday speech, but there is a real distinction”

        Mass: a fundamental measure of the quantity of matter in that body

        Weight: the weight of a body depends on the mass of the body and another body.     It is the gravitation force of one body on another

        In laboratory, masses are generally expressed in grams (g) or milligrams (mg)

        The SI (System International) unit for mass is the kilogram (kg)

Density

  • The ratio of the mass of an object to its volume

        Density → mass / volume

        Mass is usually in grams; volume in mL or cm3

Temperature

  • Related to the motion of particles that compose a substance

  • A measure of the thermal energy of a substance

            3 units of measure that are used today are: Fahrenheit, Celsius, and Kelvin

        Water freezes at 0 oC; boils at 100 oC

        Comfortable room temp → 22 oC

        Body Temp → 37 oC

        Hottest water one could put their hand into → 60 oC

                Formula

                oF=1.8oC+32^{o}F=1.8^{o}C+32

                 oC=oF−321.8^{o}C=\frac{^{o}F-32}{1.8}

Kelvin Scale

  • A temperature scale with no negative numbers

        An absolute scale:

               T(K)=oC+273T\left(K\right)=^{o}C+273

   

Scientific Notation

  • Conveniently represents any number large or small

        Product of a non-exponential (M) and exponential term (10n)

                M x 10n

            M = a number between 1 and 10; written with a decimal point to the right of the           first non-exponential term

               ex. Number                            Scientific Notation

                      180,000,000g                  1.8 × 108g

                      0.00006g                         6 × 10-5g

                      751000g                          7.51 × 105g

                      0.1590m                          1.590 × 10-1m

Chapter II - Atoms and The Periodic Table

Atomic Theory of Matter

  • The idea that matter is made of discrete units (particles)

Models of Atoms

  • Dalton Model of the Atom (John Dalton Eng. 1766 - 1844)

    • A consequence of the Law of Conservation of Matter, and the Law of Constant composition

          a. Matter is neither created nor destroyed

          b. The masses of one type of atom to another in a compound         are always the same; e.g. water is always 88.8% oxygen and         11.2% hydrogen (Joseph Proust)

    • All matter is composed of atoms—small, indivisible, indestructible particles

    • Atoms of a given element are identical

    • Different atoms have different weights and different chemical and physical properties

    • Shows that compounds are formed by combination of atoms in a ration of small whole numbers

    • Shows that a chemical r x n only involves the combination, rearrangement, or separation of atoms

  • Thompson Model of the Atom (J.J Thomson Eng. 1856 - 1909)

    • Evidence accumulates that the atoms is not indivisible. Comes from cathode ray experiment

    • The atom is composed of parts. Given credit for the discovery of the electron

    • Proposed model: a positively charged sphere with electrons placed throughout (orange is a region of diffuse positive charge)

  • Rutherford Model of the Atom (Ernest Rutherford b. NC. Br. 1871 - 1937)

    • Studied the particles given off by radioactive substances

    • In an effort to count these particles, they are passed through a thin sheet of gold metal slowing them down

    • Surprisingly deflections are detected and give a new insight into the structure of the atom

                      Interpretation of Gold Foil Experiment

      • Most particles pass through because the atom is largely empty space

      • At the center of the atom is the atomic nucleus containing positively charged particles (credited with discovering the proton)

      • Electrons are scattered about the nucleus

      • Discovers that the nucleus is very heavy and predicts that it contains a neutral particle. Later found to be the neutron (discovered by James Chadwick)

  • Bohr Model of the Atom (Neils Bohr den. 1885 - 1962)

    • Examined the light emission by different elements

    • Postulates that electrons occupy specific orbitals

    • Absorption and emission of energy is the result of electrons moving between orbital

    • Accounts for hydrogen emission lines (electrons exist in orbits at specific, fixed energies and specific fixed distances from the nucleus

  • Quantum Mechanical Model of the Atom

    • Electrons are in quantum mechanical orbitals; probability maps that show where electrons are likely to be found

Definitions for Beginning Chemists

  • Quantitative Information - usually means numerical data (e.g. the temp at which a chemical substance melts)

  • Qualitative Information - usually means non-numerical observation (e.g. the color of a substance or its physical appearance)

  • Matter - anything that has mass and occupies space; composed of different kinds of atoms; each atoms consists of a tiny core—the nucleus— each containing positive charges (protons), neutral charges (neutrons), and is surrounded by negative charges (electrons)

Elements

  • Arranged in rows (called a period) of increasing atomic number—the number of protons in the nucleus of an atom

  • Elements having similar chemical and physical properties lie in a vertical columns called a group

                    Each element has an abbreviation called a chemical symbol

                    Group A - main group (representative) elements

                    Group B - transition elements

  • Divided into several regions

    • Metals - gray

    • Non metals - green

    • Metalloids (semi-metals) - purple

Inside the Atom

  • Protons and electrons carry electric charges (positive and negative)

  • In an atom with no net electric charge, the number of negatively charged electrons around the nucleus equals the number of positively charged protons in the nucleus

    • Atomic number - the number of protons in the nucleus of an atom

    • Mass number - the sum of the masses of the protons and neutrons in the nucleus of the atom

  • Isotopes - atoms having the same atomic number but different mass numbers

    • Generally, particular isotopes are referred to by giving its mass number e.g. Uranium - 238

    • Some important isotopes have their own name

  • Atomic Mass (weight)

    • The weighted sum of the naturally occurring isotopes in a sample of the element

Electron Dot Structure (Lewis Dot Structure) of Atoms

  • Electrons orbit the nucleus in regions of space given by the laws of quantum mechanics

    • Some of these electrons are not involved in chemical reactions - core electrons - and generally do not concern us when looking reactivity

    • The other electrons - the valence electrons - are of primary importance to our understanding of reactivity

    • These electrons occupy the outermost regions of the atom

  • Symbols (electron dot or Lewis) are used to keep track these valence electrons

  • The number of valence electrons is equivalent to the group number of the element (limited to main group elements)

  • Electrons are place to maximize distance between each and are paired after four are drawn

Periodic Trends

  • Atomic size - atom radii decreases across a period and increase down a group

  • Metallic Character - lose electrons easily, shiny, malleable, electrical conducting

  • Ionization Energy - energy needed to remove an electron from a metal atom; ionizing energy increases across a period and decreases down a group