Matter - anything that has mass and occupies space

By Edith Myers,2014-04-16 07:13
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ex H2O - a water molecule, composed of 2 H atoms and 1 O atom. ethanol - C2H6O - drinkable ex salad dressing clumps of oil molecules, clumps of water, etc. IIA Alkaline earth metals ? low density, fairly reactive, form 2+ ions

Matter - anything that has mass and occupies space

Mass vs. Weight

    Mass quantity of matter

     - mass is the same everywhere, earth, moon, space

Weight depends on gravity

    - on moon .17 of the weight on earth, space 0 weight

Atoms - extremely small building blocks of matter

    all matter is composed of atoms

    atoms cannot be broken down into smaller pieces by chemical means

    Elements - the 114 known different types of atoms

     All matter is composed of the elements

    Elements cannot be decomposed into other substances

    Elements are denoted by an international chemical symbol consisting of 1 or 2


Molecule - a combination of 2 or more atoms (same or different)

    ex HO - a water molecule, composed of 2 H atoms and 1 O atom 2

    ethanol - CHO - drinkable 26

    ethylene glycol - CHO - poisonous 262slight differences in combination of atoms can have largely diff. properties

Compounds - composed of 2 or more different elements bonded together

     - can be decomposed into elements

     - ex a cube of ice composed of many HO molecules 2

ex graphite (carbon) - element or cmpd

     diamond (carbon) - element or cmpd

     carbon dioxide CO - element or cmpd 2

     oxygen Omolecule - element or cmpd 2

law of constant composition or law of definite proportions

    The chemical composition of a specific compound is always the same

    ex. there will always be exactly 2 hydrogens and 1 O in a molecule of water

    States of Matter , how many?

Matter can exist in four possible states




    Plasma the sun and lab experiments

Can all matter occupy all of these states?

    States of Matter, in the macroscopic scale table 2.1 Solid - definite shape and volume (not compressible)

    Liquid - definite volume, not shape (not compressible)

    Gas - no definite shape or volume (compressible, expandable)

    States of Matter, Submicroscopic scale Solid - particles(atoms or mol.) are packed closely in a definite arrangement

    Liquid - particles closer together than gas, particles changing location

    Gas - particles far apart, moving very quickly

State of matter depends on T, P, and nature of the substance

    intermolecular forces determine how strongly particles hold together, not


    T and P can be changed to alter phase

A Pure Substance is composed of only 1 type of atom or molecule (element or



    A Mixture is a combination of 2 or more substances, each substance retains it's own identity and it's own properties,

    - compositions of a mixture may vary (strong or weak tea)

     - concentrated or dilute

    - substances in a mixture are called components (sugar, water, etc.)

    note: air is 78% N

    , 21% O, <1% Ar, COand HO varies with humidity 222 2

    Mixtures are classified as either homogenous or heterogeneous fig 2.9

    homogenous mixture - a.k.a. solution - uniform throughout,

     ex air (mixed to molecular level)

    heterogeneous mixture - not uniform throughout

     ex salad dressing clumps of oil molecules, clumps of water, etc.

miscible, immiscible liquids

    separation of mixtures - based on retained properties of components

filtration - diff states

distillation - diff bp water 100?C, Alcohol 78?C



    Physical properties - describes the substance as is. ex. color, odor, density, mp, bp, hardness, conductivity,...

    Chemical properties - describes how the substance changes or reacts to form other substances ex. Iron will react with oxygen to form rust

    intensive properties - do not depend on the amount of the substance

     - can be used to identify a substance table 2.5

    ex. mp, density

    extensive properties - depend on amount ex mass, volume


    Physical change - change of physical appearance not composition

     ex Ice ? Water ? Water gas

     it's still HO 2being ground up, or hammered is a physical change

Names of phase changes?

    Chemical change - a reaction - alters the identity of the substance

     ex. electric current through water produces hydrogen and oxygen

    ex iron rusting - CC

    dry ice sublimes - PC

Energy and Chemical Change 2 - a rolling car Kinetic Energy The energy of motion. - a vibrating atom (thermal energy) - 1/2mv

    Potential Energy Stored energy due to position.

     - car on a hill

    - an unstable chemical structure (chemical energy)

Cars tend to roll down hills converting potential energy into kinetic energy.

    Chemical reactions tend to roll down hills too, converting chemical energy into

    thermal energy. Although not always.

Cars can be moved back to the top by adding energy.

    Many reactions can be reversed by adding energy.

    NOTE: reactions have exactly opposite energy requirements when reversed.

metabolism CHO+ 6 O ? 6 CO + 6 HO ?H = -2540 kJ 6126 (s) 2 (g)2 (g)2 (g)

    photosynthesis 6 CO + 6 HO? CHO+ 6 O ?H = 2540 kJ 2 (g)2 (g) 6126 (s) 2 (g)

    combustion of hydrogen and oxygen, and hydrolysis of water

    exothermic a rx that releases heat

    endothermic a rx that absorbs heat

    exergonic a rx that released energy

    endergonic a rx that absorbs energy

    The Law of Conservation of Energy

    a.k.a. 1st Law of Thermodynamics

    - Energy is neither created nor destroyed in chemical reactions

    (only changed in form)

Energy released or absorbed in processes always existed and always will exist in

    some form

ex car on hill to car rolling to friction produced heat

The Conversion of Matter to Energy 2

    The Einstein equation - relationship between matter and energy

     - E = mc- useful for nuclear reactions

     - proven nearly 40 years later by the atomic bomb

     - 1 g of matter to energy heat a house for 1000 years

     - less than 1% conversion in nuclear explosions

Result The sum total of matter and energy in the universe is constant.

Chapter 3 Fundamental Measurements

All measurements should include number, unit, and name of substance.

    ex 250 mg Vitamin C

    leave out anything and is senseless

3.1 Metric and SI Units

    Metric System - well organized system of units. ex cm-ml-g-cal

    conversion within unit types based on multiples of ten.

    prefixes can be used to express multiples of ten.

    the specific metric units used for scientific measurement are SI units(Systeme


    7 base units - kg - m - s - A - K - cd - mol

     - mass is kg = 2.2 pounds chemistry use gram

     - all other SI units are derived from these 7

    333- ex volume in SI is derived, m, but we usually use L(dm) or ml(cm)

Prefixes are used to indicate powers - Table 3.3 (3.2)

    Kilo k 1000 or 10


    -1 Deci d .1 or 10

    -2 Centi c .01 or 10

    -3 milli m .001 or 10

    -6micro ? .000001 or 10 -9nano n .000000001 or 10

3.7 Uncertainty in Measurement

    Precision - measurements that are in close agreement

     - measurements have more “significant figures”

     - the sloppiness of measurement is the degree of precision

    Accuracy - measurements are correct

     - How close to right the measurement is.

ex When several measurements of the same item agree with each other, they are

    precise but not necessarily accurate (perhaps the scale is wrong).

    A Thermometer marked to tenths of degrees is more precise than one marked with

    whole degrees, but not necessarily more accurate.

    If equipment is calibrated better precision leads to better accuracy.

Numbers, Exact and Inexact

    exact numbers come from defined values (3ft/yd)

     or integer count of values (14 students)

    numbers that have to be measured are always inexact (how tall? 6.????....)most

    English to metric conversions are inexact

    Significant Figures - indicates the exactness of a measurement

     ex how tall ?? how many number showing

In calculations inexact data yields inexact answers. How exact is answer?

    1st- use Rules to determine sig figs

    1. nonzero digits always sig

    ex 45 - 2 sig figs , 1.37 - 3 sig figs

    2. captive zeros must be significant

    ex 1001 - 4 sig figs , 1.0005 - 5 sig figs

    3. leading zeros are not significant

    ex .004 - 1 sig digit, 0.0045 2 sig digits

    4. trailing zeros are significant if there is a decimal point

    ex .00400 - 3 sig figs, 1000. - 4 sig figs

    5. zeros at end, no decimal point ???

    ex 1000 1,2,3,or 4 sig figs?? must assume not significant

    use exponential notation to remedy

    3ex 1000 with one sig figs is 1 ? 10

    3ex 1000 with four sig figs is 1.000 ? 10

2nd - In multiplication and division the answer should have the same number of sig

    figs as the measurement with fewest sig figs. count sig figs.

    ex 50.15/10 = 5.015 round to 1 digit, answer is 5

     .001040 * 17.295 = .0179868… round to 4 digits, answer is .01799

2nd - In addition and subtraction the result is limited by the least number of

    decimal places

    ex 20.1095

     + 1.76

     = 21.8695 round, don't truncate ? 21.87, not to 3 digits.

Dimensional Analysis - Method by which units (dimensions) are used to help

    solve problems, and check answers.

    Calculations with Dimensional Analysis 1. Use equivalence statement to get conversion factor.

    2. Pick conversion factor that cancels appropriate unit.

    3. Multiply quantity by conversion factor.

    4. Check Sig Figs.

    5. Ask whether your answer makes sense.

Conversion Factor - a fraction that expresses equivalent quantities in two different

    units, used to convert from one unit to another(like multiplying by 1)

    ex 1 inch = 2.54 cm ? conversion factor 1 in/2.54 cm or 2.54 cm/1 in

English metric conversions to know 454g = 1lb, 2.54cm = 1in, lL = 1.057qt

ex convert .3704 m to cm

     2-2 cm or .01 m = 1 cm or 10m step 1 equivalence 1m = 100 cm or 1m = 10= 1 cm

     conversion factors 1m/100cm or 100cm/1m

    step 2 - 4 .3704 m * 100 cm = 37.04 cm ? 4 sf

     1 m

    note: 1 and 100 do not limit sf

    step5 Because cm are smaller than m it makes sense that .3704 m =

    37.04 cm

ex convert 5.6 cm to inches

step 1 equivalence 2.54 cm = inch

     2 conversion factors 2.54cm/1 inch and 1inch/2.54 cm

    step 2 - 4 5.6 cm 1 inch = 2.20472 ? 2 sf = 2.2 inches

     2.54 cm

    note: 1 does not limit sf

step5 Because cm are smaller than inches it makes sense that it takes

    more cm to make inches 5.6cm = 2.2 inches

ex Convert .0478 mg to ?g

    1000mg = 1 g = 1000000 ?g or 1 mg = .001 g and 1 ?g = .000001g so 1 mg = 1000?g

.0478mg (.001g / 1mg) (1?g/.000001g)


    .0478mg (1000?g/1mg) = 47.8?g

ex convert 200. ml to fl ounces

     1 L = 1.057 qt

     1 qt = 32 fl ounces

200. ml ( 1L/1000 ml) (1.057 qt/1L) (32 fl ounces/1 qt) = 6.7648 ? 6.76 oz

    33ex convert 351 in to cm

     2.54 cm = 1 inch

    3common mistake: 351 in (2.54 cm/inch). Is made clear by units.

    3 3351 in (2.54 cm/inch) = 5751.859464 ? 5750 cc ? ?L

ex convert 95 kmph ft/s 95km/hr(1000m/km)(100 cm/m)(1in/2.54cm)(1ft/12in)(1hr/60min)(1min/60s)

     2.54 cm = 1 inch =87 ft/s

ex a faucet dripping 53 drops per minute adds how many gallons of water a month

    to the bill? 18 drops ? 1 ml

Test 1 Review

    Chapter 1

    Chemistry, chemical industry, branches chemistry, types of chemists,

    applied vs basic research

    scientific method

    Laws, Hypothesis, Theories

Chapter 2


    Mass vs. Weight

    Atoms, Elements, Molecule, Compounds

Law of constant composition or law of definite proportions

    States of Matter, in the macroscopic vs Submicroscopic scale

    Pure Substance vs mixtures

    homogenous or heterogeneous

    separation of mixtures

    Physical properties vs chemical

    intensive property vs extensive

    Physical change vs Chemical Changes

Energy and Chemical Change

    exothermic vs endothermic vs exergonic vs endergonic

The Law of Conservation of Energy

    The Conversion of Matter to Energy

Chapter 3 Fundamental Measurements

    Metric and SI Units

    Prefixes are used to indicate powers - Table 3.3 (3.2)

    3 ex 1000 N = 1 kN Kilo k 1000 or 10

    -1 ex 1 W = 10 dW Deci d .1 or 10

    -2 ex 1 L = 100 cLCenti c .01 or 10

    -3 ex 1 g = 1000mgmilli m .001 or 106-6 ex 1 m = 10 ?mmicro ? .000001 or 10 9-9 ex 1 J = 10 nJnano n .000000001 or 10

     3454g = 1lb, 2.54cm = 1in, lL = 1.057qt, 1 mL = 1 cm

    common english conversions

    12 in = 1 ft, 4 qt = 1 gal, 5280 ft = 1 mile, 16 ounces = 1 pound, 32 ounces = 1 qt

Precision vs Accuracy

    Exact and Inexact numbers

    Significant Figures

    Calculations with Dimensional Analysis

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