Lecture Presentation - Bancroft School€¦ · Matter and Measurement Based on Power Point Presentation by James F. Kirby Quinnipiac University Hamden, CT Lecture Presentation . Matter

© 2015 Pearson Education, Inc.

Chapter 1 Introduction: Matter and

Measurement Based on Power Point Presentation by

James F. Kirby Quinnipiac University

Hamden, CT

Lecture Presentation

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What is Chemistry? •  the study of the properties and behavior of matter. •  Central to our fundamental understanding of many

science-related fields.

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What is Matter? •  In general, anything that has mass and takes up space.

•  Chemical matter is composed of atoms and molecules.

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Atoms: Building Blocks of Chemical Matter •  Atoms are incredible tiny

(~10–10 m in diameter) •  Element: simplest form

of chemical matter. Each element corresponds to a unique type of atom.

•  Compound: substance

composed of two or more different of elements (i.e. two or more atom types).

Note: Balls of different colors represent atoms of different elements. Individual atoms can bond together to form larger structures called molecules.

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Classification of Matter According to State Three states of chemical matter:

Ø  Solid (e.g. ice) Ø  Liquid (e.g. water in

its standard state) Ø  Gas (e.g. water

vapor)

Click here for a YouTube tutorial video on the 3 states of matter. Be able to relate the observed properties of each state to its respective molecular-level model.

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Classifying Matter Based on Composition

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Follow this flow chart to classify any sample of matter as one of the following: Ø Element Ø Compound Ø Homogeneous

mixture Ø Heterogeneous

mixture

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Classification of Matter—Pure Substances Pure Substance: has distinct properties, and its composition does not vary from sample to sample. The two types of pure substances are elements and compounds. Ø Element: can’t be decomposed to simpler

substances. Ø Compound: can be decomposed to simpler

substances. (Composed of two or more elements.)

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Compounds and Composition •  Law of Constant Composition (or Law of Definite Proportions): A

given compound is always composed of the same two or more elements, which are always present in the same ratio by mass. (e.g. water always contains 8 grams of oxygen per gram of hydrogen.)

•  According to atomic theory, this is because the relative number of atoms of each element that makes up the compound is the same in any sample of the compound. (e.g. water is always composed of two H atoms for every one O atom, and an O atom is 16 times more massive than an H atom.)

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Mixture •  Physical blend of two or more pure

substances; components can usually be separated by physical means.

•  Exhibits the properties of the substances that make it up.

•  Heterogeneous Mixture: varies in composition throughout a sample

•  Homogeneous Mixture (or Solution): has the same composition throughout; thoroughly mixed down to the atomic/molecular level.

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Types of Properties •  Physical Property: observed without changing a

substance into another substance. ◦  Examples: boiling point, density, mass, and volume.

•  Chemical Properties: observed as a consequence of one substance changing into another substance.

◦ Examples: flammability, corrosiveness, and reactivity with acid.

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Types of Properties: Another Way to Classify Them

Extensive Properties: depend on the amount of the substance present.

Ø Examples: mass, volume, or energy.

Intensive Properties: independent of the amount of the substance that is present.

◦  Examples: density, boiling point, and color. ◦  Note: any property that is a ratio of two extensive

properties (e.g. density = mass/volume) is an intensive property.

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Types of Changes •  Physical Changes: changes in matter that do

not change the composition of a substance. ◦ Examples: changes of state, temperature, and

volume.

•  Chemical Changes: result in new substances. ◦ Examples: combustion, oxidation, and

decomposition.

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Physical Changes in State of Matter Ø Alter the physical form

but not the type of matter present

Ø Examples include changes of state (e.g. when ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.)

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Chemical Reactions (Chemical Changes) Ø The starting substances

(reactants) are converted to new substances (products) with different properties.

Ø Here, the elements hydrogen (H2) and oxygen (O2) become the compound water (H2O).

Click here to view a video that demonstrates the differences between chemical and physical properties and changes. Also covered are the topics of elements, compounds, and mixtures.

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Separating Mixtures

Mixtures can be separated based on physical properties of the components of the mixture. Some methods used to separate the components of mixtures are

Ø Filtration Ø Distillation Ø Chromatography

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Filtration •  Solid substances

are separated from liquids and solutions

•  Liquid portion passes through pores in the filter paper; while solid particles cannot.

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Distillation •  Uses differences in the

boiling points of substances to separate a homogeneous mixture into its components.

•  Relative to the boiling liquid mixture, the vapor is more concentrated in the more volatile component

•  The vapor is then condensed and collected in a “receiving flask.”

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Chromatography •  A mixture is dissolved in

a solvent (mobile phase) that carries it through the chromatography column.

•  Substances separate based on differences in their abilities to adhere to the solid surface (stationary phase)

•  In this case, “substance a” adheres most strongly to the stationary phase and, thus, comes out last. “Substance c” adheres least strongly and comes out first.

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Numbers and Chemistry •  Chemistry is very quantitative, as it often

involves things with numerical values (i.e. experimental measurements).

•  Important Numerical Concepts in Science Ø Scientific Notation Ø Units of measurement Ø Quantities that are measured and calculated Ø Uncertainty in measurement Ø Significant figures Ø Dimensional analysis

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Units of Measurement: SI Units

•  Système International d’Unités (International System of Units) •  Each fundamental quantity has its own base (see above table). •  Derived SI units are constructed from the appropriate base units

(e.g. the SI unit for velocity is meters per second, m/s).

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Scientific Notation Convenient way of expressing both the very large and very small quantities often encountered in science

o Example: There are 6 x 1023 atoms in 1 gram of hydrogen. o Example: The mass of a proton is 1.67 x 10–27 kg.

Click here for a YouTube video tutorial on Scientific Notation. This 14 minute video gives lots of examples.

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Metric Prefixes: convert base units into units that are more common and/or convenient

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Mass and Length

•  Two of the most basic units commonly measured in science.

•  Mass: measure of the amount of matter in an object. SI unit is the kilogram (kg).

•  Length: measure of distance. SI is the meter (m).

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Volume: Example of a Derived Unit •  Derived from a base/fundamental unit:

length •  The SI unit of volume is the cubic meter:

m × m × m = m3. •  More commonly used metric units of

volume: liter (L) and milliliter (mL). Ø A liter is a cube 1 decimeter (dm) long

on each side. (1 L = 1 dm3) Ø A milliliter is a cube 1 centimeter (cm)

long on each side, also called 1 cubic centimeter: cm × cm × cm = cm3. (1 mL = 1 cm3)

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Temperature v Generally related the “hotness and coldness” of an object. v Determines the direction of heat flow. Heat always flows

spontaneously from a higher temperature to a lower temperature. v Higher temperatures correspond to more thermal energy (random

atomic/molecular motion) in a substance. v  In scientific measurements, the Celsius and Kelvin scales are used

most often used.

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Temperature Units Used in Science Celsius scale: based on the properties of water.

Ø  0 ˚C is the freezing point of water. Ø  100 ˚C is the boiling point of water.

Kelvin scale: the “absolute” temperature Ø  The SI unit of temperature. Ø  The Kelvin unit has the same size as the Celsius unit, but the

two scales have different zero points Ø  0 K is the lowest possible temperature (“absolute zero”). This

is the point at which all thermal energy would have been drained from an object; so the temperature cannot get any lower.

Ø  0 K = –273.15 ˚C

Conversion Between Scales (given on AP Exam) K = ˚C + 273.15

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Density o  The mass to volume ratio of a substance:

D = m/V o  a physical property, with units derived from units for

mass and volume. •  The most common density unit for liquids and solids is

g/mL (Note: 1 g/ml = 1 g/cm3) For example, the density of aluminum at 20 ˚C is 2.7 g/cm3.

•  A common density unit for gases is g/L. For example, the density of air at sea level and 15 ˚C is 1.2 g/L.

•  The SI unit of density is kg/m3, but it is not commonly used. For example, the density of aluminum at 20 ˚C is 2.7 x 103 kg/m3.

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Types Numbers Encountered in Science Exact Numbers Ø  Either counted explicitly (e.g. there are 16 students in a

classroom) or given by definition (e.g. there are 12 eggs in 1 dozen)

Ø  There is no uncertainty in their values. Inexact (Measured) Numbers: Ø  Involve a certain level of estimation, which is dependent

upon the circumstances of each measurement Ø  Uncertainty is inherent in any measurement, since both

scientists and scientific instruments have limitations. (For example, some scientists have better eyesight than others. Also, some balances measure to ±0.01 g; others measure to ±0.0001g.)

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Accuracy versus Precision Accuracy ²  refers to the closeness of a measurement

to the true value of a quantity. ² Based on comparing the measured value

to the true/accepted value for the quantity.

Precision ²  refers to the proximity of several

measurements to each other. ² Based on comparing the measured values

to each other. (The true value is not taken into account when determining precision.)

Click here to view a tutorial video on the difference between accuracy and precision.

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Significant Figures

•  Each measured or calculated value has a certain number of significant figures. These are the digits in which there is a good level of confidence.

•  Measuring the smallest significant digit involves

a certain level of estimation. This determines the precision (i.e. level of reproducibility) of the measured value.

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Uncertainty in Measurements •  Each measurement involves a certain degree of inherent random error,

which limits the precision (reproducibility) of the measured value. •  The amount of precision (i.e. number of sig figs) in a measurement is

largely determined by the instrument used. For instance, the graduated cylinder shown below measures volume to a precision of ± 1 mL, while the burette is precise to ± 0.1 mL.

Click here to view a video tutorial on determining the numbers of significant figures in measured values.

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Rules for Determining the Number of Significant Figures in a Numerical Value

1.  All nonzero digits are significant. 2.  Zeroes between two significant figures are

themselves significant. 3.  Zeroes at the beginning of a number are never

significant. 4.  Zeroes at the end of a number are significant if

a decimal point is written in the number.

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Rules for Determining the Number of Significant Figures in a Calculated Value

Note: the precision of a calculated value is limited by the precision of the numbers used in the calculation. We must round the calculated value to the correct number of significant figures so as to not overstate the precision of the answer. •  For addition or subtraction, answers are rounded to the least

significant decimal place. •  For multiplication or division, answers are rounded to the number

of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

Click here to view a video with a comprehensive review of how to determine the numbers of sig figs in both given numbers and calculated values.

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Dimensional Analysis •  Used to convert one quantity to another. •  Often utilizes conversion factors between equivalent quantities

(e.g., 1 in. = 2.54 cm exactly, by definition). •  For instance, the equivalence between inches and centimeters can

be used to convert from cm to in. (1 in/2.54 cm) or from inches to cm (2.54 cm/1 in.) We use the ratio which allows us to change units (puts the units we have in the denominator to cancel).

Click here to view a video tutorial involving a simple one-step unit conversion of a measured value.

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Additional Video Tutorials from YouTube

Click here to view a tutorial video that includes two multi-step conversion problems.

Click here to view a video covering a numerical word problem involving a density calculation.

Click here to view a video covering a numerical word problem to determine the diameter of a red blood cell.