Chapter Outline

Chapter 1: Matter and Measurement 1.1 Elements and Atoms

Elements – composed of only one type of atom 113 elements known, ~90 found in nature Many elements have names and symbols with Latin or Greek elements Periodic Table- includes the symbol and other information about the elements Atom- smallest particle of an element that retains the characteristic of that

element. (A)Na – Sodium, Cl – Chlorine, Cr – Chromium (B)Zinc – Zn, Nickel – Ni, Potassium-K

1.2 Compounds and Molecules Chemical Compound- A pure substance with is composed of two or more

different elements ~20 million compounds now known When elements become part of a compound, their original properties (Ex.

Color, hardness, melting point) are replaced by the characteristic properties of the compound.

Compounds have distinctly different characteristics from their parent elements and have a definite percentage composition (by mass) of their combining element.

Ions- electrically charged atoms or groups of atoms Molecules- Smallest, discrete units that retain the composition and chemical

characteristics of the compound. Chemical Formula- Any compound can be represented by it. (ex. H20)

(a) Iron – solid, cube, gray; (b) Water- liquid, colorless; (c) table salt- colored, white, solid, (d) colorless, odorless, gaseous1.3 Physical Properties- can be observed and measured without changing the composition

of a substance (ex. Color, state of matter, melting point, boiling point, density, solubility, electric conductivity, malleability, ductility, viscosity) Density- the ration of the mass of an object to its volume, is a physical property

useful for identifying substances. Density = Mass / Volume

o The density of a substance relates the mass and volume of a substance. If any two of three quantities – mass, volume, and density – are known for a sample of matter, the third can be calculated.

o Mass (g) = volume x density = volume (cm3) x Mass (g) / Volume (cm3) Temperature- the property of matter that determines whether heat energy can be

transferred from one body to another and the direction of that transfero Three temperature measurement scales are commonly used: Fahrenheit,

Celsius (used for measurements in the laboratory), and Kelvin scales (When calculations incorporate temperature data).

o Celsius- used in most other countries and in scientific notation- 0ºC = freezing point of water, 100ºC = Boiling Point

o Kelvin scale- international standard for science, uses same unit as Celsius scale but uses lowest possible temp. as its zero, a point called absolute zero. (Freezing point of water- 273.15 K; Boiling Point- 373.15 K)

o Celsius converted into Kelvins using T(K) = 1 K / 1 ºC(T ºC + 273.15 ºC) Temperature Dependence of Physical Properties

o The temp. of a sample matter often affects the numerical values of its properties

o Because the density of liquids changes with temperature, it is necessary to report the temperature when you make accurate volume measurements. (Lab glassware specifies the temp. at which they were calibrated)

Extensive and Intensive Propertieso Extensive properties- depend on the amount of substance present.

(Ex. mass and volume)o Intensive Propereties- Do not depend on the amount of substance

(ex. Ice will melt at 0º C whether an ice cube or iceberg)

1.4 Physical and Chemical Change Physical change- changes in physical properties- identity of substance is

preserved Chemical change/reaction- one of more substances (reactants) has been

transformed into one or more different substances (products) Chemical Property- involves a change in the identity of a substance A chemical change produces a new arrangement of atoms without a gain or

loss in the number of atoms of each kind The particles (atoms, molecules, or ions) present after the reaction, however,

are different from those present before the reaction. Chemical Equation- the representation of the change with chemical formulas. A physical change does not result in a new chemical substance. The substances (atoms, molecules, or ions) present before and after the change

are the same, but their arrangement relative to one another is different. Physical and Chemical changes are often accompanied by a transfer of energy.

Water to Steam – Liquid to gas (physical change) Wood Burning (chemical change)1.5 Classifying Matter

States of Matter and the Kinetic-Molecular Theoryo State of Matter- Whether a substance is a solid, liquid, or gas.

Solid- Rigid shape, fixed volume that changes little as temp and pressure change.

Liquids- fixed volume, but takes shape of its container and has no definite shape.

Gases- volume of gas is volume of container, volume of a given amount of gas varies with temperature and pressure.

o Kinetic-Molecular Theory- help us interpret the properties of solids, liquids, and gases.

States all matter consists of extremely tiny particles(atoms, molecules, or ions) which are all in constant motion

In solids these particles are packed closely together usually in a regular array.

The particles vibrate back and fourth and seldom squeeze past its neighbors.

In liquids or gases atoms or molecules are arranged randomly. Are fluid because the particles are not confined to specific

locations and can move past one another. Gas particles are far apart and move extremely rapidly and

randomly, allowing them to fill their containers. The higher the temp. the faster the particles move. The particles energy of motion (kinetic energy) acts to overcome

the forces of attraction between the particles. A solid melts to form a liquid when the temp. of the solid is raised

to the point at which the particles vibrate fast enough and far enough to push one another out of the way and move out of their regularly spaced positions.

As temp increases even more, the particles move even faster until finally they can escape and enter the gaseous state.

Matter at the Macroscopic and Particulate Levelso Observations and manipulation generally take place in the macroscopic

world of chemistry.o Submicroscopic/Particulate- the level of individual particles that make up

all of matter. Pure Substances

o Every pure substance has two features: It has a set of unique properties by which it can be recognized It cannot be separated into two or more different pure substances

by and physical technique Mixtures: Homogeneous and Heterogeneous

o Heterogeneous mixture- a mixture in which the uneven texture of the material can be detected.

o Homogeneous mixture- completely uniform at the particulate level and consists of one or more substances in the same phase.

o Solutions- Composition is the same everywhere in the sample. (ex. Air)o Purified- a mixture is separated into its pure components.(repetition

increases the purity)1.6 Units of Measurement

Qualitative observations- Color or appearance of the substances; no measurements involved.

Quantitative observations- Involve numerical observations. (time, mass, volume, and distance)

Mass- Kilogram (Kg), Length- meter (m), Time- second (s), Temperature- Kelvin (K), Amount of Substance- mole (mol), Electric Current- Ampere (A)

1.7 Using Numerical Information Length Volume

o Handled in glassware such as beakers, flasks, pipets, graduated cyclinders, and burets, which are marked in volume units.

o SI unit of voume is the cubic meter (c3), which is too large for everyday laboratory use.

o Chemists usuallu use liter (L) instead, A cube with sides equal to 10 cm (0.1 m) has a volume of 10cm x 10cm x 10cm = 1000 cm3

o 1 liter (L) = 1000 mL = 1000 cm3

o 1 cm3 = 0.001 L = 1 milliliter (1 mL) Mass

o The mass is the fundamental measure of the quantity of matter in the body and the SI unit of mass is the kilogram (kg) Smaller masses use grams (g) or milligrams (mg) 1 kg = 1000 g, 1 g = 1000 mg

Making Measurements: Precision, Accuracy, And Experimental Erroro Precision- indicates how well several determinations of the same quantity

agree. o Accuracy- the agreement of a measurement with the accepted value of the

quantity. o Average Deviation- precisions of a measurement- calculate the difference

between each experimental result and the average result. These differences, each expressed as a positive quantity, are averaged, and the results of the experiment are reported as the average value plus or minus the average deviation.

o Percent Error- the difference between your result and the accepted value; Error = experimentally determined value – accepted valuePercent Error = Error in measurement / Accepted value x 100%

Significant Figures