Chm 260 notes: phases and classification of matter (lecture 1)

The classification of matter

• been established that matter is made up of atoms

• matter: anything that occupies space and has mass

• everything is made up of matter

The state of matter

• we can classify matter according to its state (physical form) and composition (basic components that make it up)

• matter can be a solid, liquid, or gas based on what properties it exhibits

• state of matter changes from solid to liquid to gas with an increasing temperature

Structure determines properties

atoms or molecules have different structures (arrangements) in solids, liquids, and gases leading to different properties

Solid

• in solid matter, atoms or molecules pack close to each other in fixed locations

• a solid has a fixed volume and rigid shape (ice, aluminum, diamonds)

Crystalline and amorphous solid matter

• solid mater may be crystalline so atoms or molecules are in patterns with a long range and repeating order (table salt and diamonds)

• solid matter may also be amorphous so atoms or molecules do not have any long range order (glass)

Liquid

• atoms or molecules pack about as closely as they do in solid matter but they are free to move relative to each other

• liquids have a fixed volume but not a fixed shape

• liquids ability to flow makes them assume the shape of their container (examples are water, alcohol, and gasoline which are liquids at room temperature)

Gas

• atoms or molecules have a lot of space between them

• they are free to mov relative to one another

• these qualities make gases compressible and have no fixed volume or shape

Comparing solid, liquid, and gas matter

• solid matter: atoms are close packed and are not free to move

• liquid matter: atoms are close packed and atoms are free to move

• gas matter: atoms are not close packed and atoms are free to move

The classification of matter by components

• pure substance: made up of only one component and its composition is invariant

• mixture: a substance composed of two or more components in proportions that can vary from one sample to another

Classification of pure substances

• pure substance are categorized into elements and compounds which depends on whether or not they can be broken down (or decomposed) into simpler substances

Element vs. a compound

• element: a substance that cannot be chemically broken down into simpler substances (basic building blocks of matter and are composed of a single type of atom, like helium)

• compound: substance composed of two or more elements in fixed definite proportions

• most elements are chemically reactive and combine with other elements to form compounds like water, sugar, etc.

Classification of mixtures

• mixtures can either be heterogenous or homogenous

• this categorization of mixture depends on how uniformly the substances within them mix

Heterogeneous mixture

• one in which the composition varies from one region of the mixture to another

• made of multiple substances whose presence can be seen (salad or salt and sand mixture)

• portions of a sample of heterogenous mixture have different composition and properties

Homogeneous mixture

• one made of multiple substances but appears to be one substance

• all portions of a sample have the same composition and properties (like sweetened tea)

• homogeneous mixtures have uniform compositions because the atoms or molecules that compose them mix uniformly

Separating mixtures

• mixtures are separable because the different components have different physical or chemical properties

• various techniques that exploit these differences are used to achieve separation

• a mixture of sand and water can be separated by decanting- carefully pouring off the water into another container

Separating a homogeneous mixture

• can be usually separated by distillation, a process in which the mixture is heated to boil off the more volatile (easily vaporizable) liquid. The volatile liquid is then re-condensed in a condenser and collected in a separate flask

Filtration

a mixture of an insoluble solid and a liquid can be separated by filtration, which is a process where the mixture is poured through the filter paper in a funnel

Physical changes

• changes that alter only the state or appearance but not composition

• the atoms or molecules that compose a substance do not change their identity during a physical change

• when water boils it changes its state from a liquid to a gas

• the gas remains composed of water molecules (physical change)

Chemical changes

• changes that alter the composition of matter

• during a chemical change, atoms rearrange, transforming the original substances into different substances

• rusting of iron is a chemical change

Physical property

• property that a substance displays without changing its composition

• the smell of gasoline

• odor, taste, color, appearance, melting point, boiling point, and density

Chemical property

• property a substance displays only by changing its composition through a chemical change (or chemical reaction)

• the flammability of gasoline, in contrast, is a chemical property

• chemical properties include corrosiveness, acidity, and toxicity

Intensive and extensive property

• intensive property: does not depend on the amount of matter present (ex. temperature)

• extensive property: depends on the amount of matter present (mass and volume)

The units of measurement

• in chemistry, units- standard quantities used to specify measurements- are critical

• the two most common unit systems are as follow: metric and imperial system

• scientists use the international system of units (SI) which is based on the metric system

The Meter: A Measure of Length

• the meter (m) is slightly longer than a yard (1 yard is 36 inches while 1 meter is 39.37 inches)

• 1 meter=1/10000000 of the distance from the equator to the North Pole (through Paris)

Kilogram: a Measure of Mass

• the mass of an object is a measure of the quantity of matter within it

• the SI unit of mass= kilogram (kg)

• 1 kg= 2 lb 3 0z

• a second common unit of mass is the gram (g)

• one gram=1/1000kg

• weight of an object is a measure of the gravitational pull on its matter

The Second: A Measure of Time

• measure of the duration of an event

• SI units=second (s)

• other units: minute (min)=60s; hour (hr)=60 min. So 1 hr=360 seconds

The Kelvin: A Measure of Temperature

• the kelvin (K) is the SI unit of temperature

• the temperature is a measure of the average amount of kinetic energy of the atoms or molecules that compose the matter- the hotter a material the more energetic (more active) the molecules that make up the material

• Temperature also determines the direction of thermal energy transfer or what we commonly call heat

• Kelvin scale (absolute scale) assigns 0K (absolute zero) to the coldest temperature possible

• absolute zero (-273 C or -459 F) is the temperature which molecular motion virtually stops. Lower temperatures do not exist

Kevin: A Measure of Temperature equation

• the fahrenheit degree is 5/9ths the size of a celsius degree

• the celsius degree and kelvin degree are the same size

• K=C+273.15 or just 273

• C=(F-32)/1.8 (C is celsius)

Prefix multipliers

• multipliers that change the value of the unit by the powers of 10 (just like an exponent does in scientific notation)

• example: kilometer has kilo in the front so its equivalent to 1000 or 10³

Actual prefix multipliers that need to be memorized

• kilo=k=1000/10³

• milli=m=.001/10^-3

• micro=u=.000001/10^-6

• nano=n=.000000001/10^-9

• pico=p=.000000000001/10^-12

Counting Significant figures

the greater the number of sig fids, the greater the certainty of the measurement

Sig fig rules

1. all nonzero digits are significant

2. interior zeroes (zeroes between two nonzero digits) are significant

3. leasing zeroes (zeroes to the left of the first nonzero digit) are not significant. they only serve to locate the decimal point

4. trailing zeroes (zeroes at the end of a number) are categorized as follows:

• trailing zeroes after a decimal point are always significant

• trailing zeroes before a decimal point (and after a nonzero number) are always significant

Exact numbers

• have an unlimited number of sigfigs

• exact counting of discrete objects

• integral numbers that are part of an equation

• defined quantities

• some conversion factors are defined quantities while others are not

Sigfigs in calculations

we should not lose or gain precision during mathematical operations

you can not be more precise than your least precise number

Rules for calculations (multiplication and division)

the results carries the same number of significant figures as the factor with the fewest significant figures

Rules for adding and subtracting

the result carries the same number of decimal places as the quantity with the fewest decimal places

Digit to be dropped rule

if the digit to be dropped (the one immediately to the right of the digit to be retained) is less than 5, “round down” and leave the retained digit unchanged; if it is more than 5, “round up” and increase the retained digit by 1. If the dropped digit is 5 and its either the last digit in the number or its followed only by zeros, round up or down, whichever yields an even value for the retained digit. If any nonzero digits follows the dropped 5, round up.

Examples of sigfigs

• 15.51 rounds to 16

• 15.5 rounds to 16

• 15.500001 rounds to 16

• 12.500000 rounds to 12

• 15.49 rounds to 15.4

Rounding in multistep calculations

• avoid rounding errors in multistep calculations round only in the final answer

• do not round intermediate steps, if you write down intermediate answers, keep track of sig figs by underlining the least significant digit

Precision and accuracy

• accuracy refers to how close the measured value is to the actual value

• precision refers to how close a series of measurements are to one another or how reproducible they are