Slide 2 Matter and Measurement Chapter 1 Slide 3 What Is Matter?
Matter is anything that takes up space and has mass. Mass is the
amount of matter in an object. Mass is resistance to change in
motion along a smooth and level surface. Slide 4 Types of Matter
Substance- a particular kind of matter pure Fixed composition
Distinct properties Ex. Water, Salt Mixture- more than one kind of
matter Compositions vary Each substance retains its own chemical
identity and properties. Slide 5 Substances Elements- simplest kind
of matter Cannot be broken down into simpler All one kind of atom.
Compounds- are substances that can be broken down by chemical
methods When they are broken down, the pieces have completely
different properties than the compound. Made of molecules- two or
more atoms Slide 6 Mixtures Heterogeneous- mixture is not the same
from place to place. Chocolate chip cookie, gravel, soil. Variable
Composition Homogeneous- same composition throughout. Kool-aid,
air. Every part keeps its properties. Slide 7 Solutions Homogeneous
mixture Mixed molecule by molecule Can occur between any state of
matter. Solid in liquid- Kool-aid Liquid in liquid- antifreeze Gas
in gas- air Solid in solid - brass Liquid in gas- water vapor Slide
8 Solutions Like all mixtures, they keep the properties of the
components. Can be separated by physical means Not easily
separated- can be separated. Slide 9 Gold 24 karat gold 18 karat
gold 14 karat gold Gold Copper Silver 18 / 24 atoms Au 24 / 24
atoms Au 14 / 24 atoms Au Slide 10 Solid Brass An alloy is a
mixture of metals. Brass = Copper + Zinc Solid brass homogeneous
mixture a substitutional alloy Copper Zinc Slide 11 Brass Plated
Brass = Copper + Zinc Brass plated heterogeneous mixture Only brass
on outside Copper Zinc Slide 12 Steel Alloys Stainless steel
Tungsten hardened steel Vanadium steel We can engineer properties
Add carbon to increase strength Too much carbon too brittle and
snaps Too little carbon too ductile and iron bends Slide 13 Nitinol
Wire Alloy of nickel and titanium Remembers shape when heated
Applications: surgery, shirts that do not need to be ironed. Slide
14 Properties Physical Properties- A change that changes
appearances, without changing the composition, thus identity is
preserved. ex. color, odor, density. Chemical Properties- a
property that can only be observed by changing the type of
substance. ex. flammability Slide 15 Physical and Chemical Changes
Chemical Changes Physical Changes Rusting NailMelting Ice Bleaching
a StainBoiling Water Burning a LogSawing a log in half Tarnishing
SilverTearing Paper Fermenting of Grapes Breaking a Glass Souring
of MilkPouring of Milk Slide 16 Intensive Vs. Extensive Properties
Intensive Properties Do not depend on on the amount of sample being
examined Aid in the identification of a substance. ex. temperature,
density, melting point. Extensive Properties Depend on the quantity
of sample ex. mass, volume, area. Slide 17 The SI System The SI
system has seven base units from which all others are derived
Physical Quantities Name of UnitAbbreviation MassKilogramKg
LengthMeterM TimeSecondS Electric CurrentAmpereA TemperatureKelvinK
Luminous IntensityCandelaCd Amount of Substance Molemol Slide 18 SI
Units (Cont) These prefixes indicate decimal fractions or multiples
of various units PrefixAbbreviationMeaning Mega-M10 6 Kilo-k10 3
Deci-d10 -1 Centi-c10 -2 Milli-m10 -3 Micro- 10 -6 Nano-n10 -9
Pico-p10 -12 Femto-f10 -15 Slide 19 Temperature Conversions At home
you like to keep the thermostat at 72 F. While traveling in Canada,
you find the room thermostat calibrated in degrees Celsius. To what
Celsius temperature would you need to set the thermostat to get the
same temperature you enjoy at home ? C = 5/9 ( F-32) F = 9/5 (C )
+32 K = C + 273.15 Slide 20 Derived Units SI units are used to
derive the units of other quantities. These units express speed,
velocity, area and volume. They are either base units squared or
cubed, or they define different base units Slide 21 Volume
Calculated by multiplying L x W x H Basic SI unit of volume is the
cubic meter (m 3 ). Smaller units are sometimes employed ex. cm 3,
dm 3 . Volume is more commonly defined by liter (L). Slide 22
Density Density is an intensive property of all substances; this
Means that it is a universal characteristic of the substance and
does not change because of extensive or accidental properties such
as amount, size or location. Density is the relationship between
the mass and volume of an object. It is defined as: Density = Mass
of substance Volume of substance Density is usually expressed in
g/ml Slide 23 Specific Gravity Specific Gravity is a ratio between
the density of a substance and the density of water. It is defined
by: Specific Gravity = Density of Sample Density of Water In
calculations involving s.g., the units of density must match in
order for the units to cancel Slide 24 Problems Involving Density
1. The mass of 325 mL of the liquid methanol is found to be 257 g.
What is the density of methanol? 2. A lead weight used in the belt
of a scuba diver has a mass of 226g. When the weight is carefully
placed in a graduated cylinder containing 200.0ml of water, the
water level rises to 220.0ml. What is the density of the lead
weight (g/ml) 3.A sample of a certain material has a mass of 2.03x
10 -3 g. Calculate the Volume of the sample given that the density
is 9.133 x 10 -1 g/cm 3 Slide 25 Uncertainty and Measurement Exact
numbers Numbers whose values are known exactly e.x. 12 eggs in a
dozen, 1000g in a kg Inexact numbers Numbers obtained by measuring
a quantity e.x. height, weight or temperature There is always a
degree of uncertainty in measured values. Why? Slide 26 Precision
Vs Accuracy When comparing sets of data points, scientist want to
know two things: which set of data points are precise and which set
are accurate. Slide 27 Precision How many times a given measurement
can be repeated with results close in value to each other. Which of
the following data is more precise: 119g, 120g, 128g or 101g, 100g,
99g Good precision poor accuracy Slide 28 Accuracy How close an
experimental measurement is to the true, actual, or book value
Usually the more accurate a measurement the more precise it will
be. Good Accuracy and Precision Slide 29 11/29/201528 Precision and
Accuracy Accuracy refers to the agreement of a particular value
with the true value. Precision refers to the degree of agreement
among several measurements made in the same manner. Neither
accurate nor precise Precise but not accurate Precise AND accurate
Slide 30 Significant Figures When using a measuring device to
measure anything the last number of the measurement is always
estimated Measured quantities are usually reported in such a way
that only the last digit is uncertain. All digits including the
uncertain digit are called significant figures Slide 31 Significant
Figures A number is not significant if it is: A zero at the
beginning of a decimal number ex. 0.0004lb, 0.075m A zero used as a
placeholder in a number without a decimal point ex. 992,000,or
450,000,000 A number is a S.F. if it is: Any real number ( 1 thru
9) A zero between nonzero digits ex. 2002g or 1.809g A zero at the
end of a number or decimal point ex. 602.00ml or 0.0400g Any digit
in the coefficient of a number written in scientific notation ex.
4.0 x 10 5 m, 5.70 x 10 -3 Slide 32 11/29/201531 Sig Fig Practice
#1 How many significant figures in each of the following? 1.0070 m
5 sig figs 17.10 kg 4 sig figs 100,890 L 5 sig figs 3.29 x 10 3 s 3
sig figs 0.0054 cm 2 sig figs 3,200,000 2 sig figs Slide 33
Addition and Subtraction When adding or subtracting the answer is
rounded so it has the same number of decimal places as the number
with the least number of decimal places 6.23 cm 39.24 cm + 677.1 cm
722.6 cm Slide 34 Multiplication and Division When multiplying or
dividing the answer has the same number of significant digits as
the number with the least number of S.F. 2.85ml x 67.4ml = 192ml ml
49.618g = 43.8ml Slide 35 11/29/201534 Sig Fig Practice #2 3.24 m +
7.0 m CalculationCalculator says:Answer 10.24 m 10.2 m 100.0 g -
23.73 g 76.27 g 76.3 g 0.02 cm + 2.371 cm 2.391 cm 2.39 cm 713.1 L
- 3.872 L 709.228 L709.2 L 1818.2 lb + 3.37 lb1821.57 lb 1821.6 lb
2.030 mL - 1.870 mL 0.16 mL 0.160 mL Slide 36 Odd Numbers Odd
numbers are rounded up, even numbers are left alone when the
remainder is 5 ex. 22.15 g if rounded to 3 SF = 22.2g 22.25 g if
rounded to 3 SF = 22.2 Slide 37 11/29/201536 In science, we deal
with some very LARGE numbers: 1 mole = 602000000000000000000000 In
science, we deal with some very SMALL numbers: Mass of an electron
= 0.000000000000000000000000000000091 kg Scientific Notation Slide
38 11/29/201537 2 500 000 000 Step #1: Insert an understood decimal
point. Step #2: Decide where the decimal must end up so that one
number is to its left up so that one number is to its left Step #3:
Count how many places you bounce the decimal point the decimal
point 1234567 8 9 Step #4: Re-write in the form M x 10 n Slide 39
11/29/201538 2.5 x 10 9 The exponent is the number of places we
moved the decimal. Slide 40 11/29/201539 0.0000579 Step #2: Decide
where the decimal must end up so that one number is to its left up
so that one number is to its left Step #3: Count how many places
you bounce the decimal point the decimal point Step #4: Re-write in
the form M x 10 n 12345 Slide 41 11/29/201540 5.79 x 10 -5 The
exponent is negative because the number we started with was less
than 1. Slide 42 11/29/201541 Review: Scientific notation expresses
a number in the form: M x 10 n 1 M 10 n is an integer Slide 43
Conversion Factors Fractions in which the numerator and denominator
are EQUAL quantities expressed in different units Example: 1 in. =
2.54 cm Factors: 1 in. and 2.54 cm 2.54 cm 1 in. Slide 44 Learning
Check Write conversion factors that relate each of the following
pairs of units: 1. Liters and mL 2. Hours and minutes 3. Meters and
kilometers Slide 45 How many minutes are in 2.5 hours ? Conversion
factor 2.5 hr x 60 min = 150 min 1 hr 1 hr cancel By using
dimensional analysis / factor-label method, the UNITS ensure that
you have the conversion right side up, and the UNITS are calculated
as well as the numbers! Slide 46 Dimensional Analysis Many problems
in chemistry and health sciences require a change in units, thus
dimensional analysis is used to make these conversions The key is
the correct use of a conversion factor to change 1 unit to another
ex. 1hr = 60 min this allows us to write the relationship: 60 min
and 1hr 1hr60min Slide 47 Problems Involving Dimensional Analysis
1.On a bicycle trip, Maria averaged 35 miles per day. How many days
did it take her to cover 175 miles? 2.To prevent bacterial
infection, a doctor orders 4 tablets of amoxicillin per day for 10
days. If each tablet contains 250 mg of amoxicillin, how many
ounces of the medication are given in 10 days? Slide 48 Solution to
Problems Step 1: Given 175 miles 35 miles = 1 day Step 2: Unit plan
Miles Days 175 miles x 1 day = 5.0 days 35 miles Slide 49 Solution
to Problem Step 1: Given 10 days 4 tablets/day 250 mg each Step 2:
Unit Plan mg g lb oz 1000mg x 1g x 1lb x 16 oz = 0.35 oz 1000mg
453.59g 1lb Slide 50 Dealing with Two Units If your pace on a
treadmill is 65 meters per minute, how many seconds will it take
for you to walk a distance of 8450 feet? Slide 51 Learning Check A
Nalgene water bottle holds 1000 cm 3 of dihydrogen monoxide (DHMO).
How many cubic decimeters is that? A Nalgene water bottle holds
1000 cm 3 of dihydrogen monoxide (DHMO). How many cubic decimeters
is that? Slide 52 Solution 1000 cm 3 1 dm 3 10 cm 10 cm ( ) = 1 dm
3 So, a dm 3 is the same as a Liter ! A cm 3 is the same as a
milliliter.
