Introduction to Chem I

By Luis Bello

Introduction

Why studying chemistry is important:

1. Understanding the World: Chemistry explains matter’s composition, properties, and reactions, helping us make sense of the natural world and everyday phenomena.

2. Foundational Science: Chemistry’s concepts underpin other sciences like biology, physics, and geology, linking multiple fields of study.

3. Practical Applications: Chemistry drives innovations in medicine, materials, food, fuels, and cleaning products, addressing practical needs and challenges.

Why studying chemistry is important:

1. Preparing for STEM Careers: Chemistry is crucial for careers in medicine, engineering, science, and technology, providing a strong foundation for various professional opportunities.

2. Developing Scientific Thinking: Chemistry enhances skills in data gathering, problem solving, and analysis, which are valuable in STEM fields and everyday life.

3. Appreciation for the Scientific Method: Chemistry illustrates the process of experimentation, data analysis, and hypothesis testing, fostering a deeper understanding and appreciation of the scientific method.

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CHEM I introduces fundamental chemistry concepts, including atomic structure, chemical bonding, molecular geometry, periodic trends, stoichiometry, chemical reactions, kinetics, equilibrium, acids and bases, and thermodynamics.

The course features lectures and hands-on laboratory sessions with chemical techniques and instrumentation.

Evaluation includes exams, lab reports, and problem sets.

CHEM I provides essential foundational knowledge for further study in chemistry and related fields such as engineering, medicine, and environmental science.

About me

• Bachelor on Chemical Education (1982). Universidad de Oriente, CUBA

• ​Dr on Chemical Education (1993), Universidad de Oriente, Cuba

• ​+ 30 years teaching Chemistry (General Chemistry, Inorganic Chemistry, Algebra, Spanish).

• Mother Language Spanish

• ​Like to travel a lot.

• ​4 Children, married,

• ​Like to work with computers.

• Like to ride bikes and going to the beach.

Textbook

Grading

Lab activities​

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Introduction of each other.

Today

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Textbook

• ​TextBook: Chemistry Atoms First 2e (OpenStax):

Atoms First Textbook for Download - OpenStax

• ​Other Resources I use:

1. ​https://courses.lumenlearning.com/trident-boundless-chemistry/

2. ​https://opentextbc.ca/introductorychemistry/

3. ​https://youtu.be/ZuWa827qAao

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Open Ended

What is Chemistry?

Chemistry is the study of matter and the chemical reactions between substances. Chemistry is also the study of matter’s composition, structure, and properties. Matter is essentially anything in the world that takes up space and has mass.

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Chemistry is sometimes called “the central science” because it bridges physics with other natural sciences, such as geology and biology.

• Sub-domains of chemistry include: analytical chemistry, biochemistry, inorganic chemistry, organic chemistry, physical chemistry, and biophysical chemistry.

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What is Chemistry?

Chemistry: The Central Science

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• matter: Something that has mass and takes up space (has volume) and makes up almost everything in the world.

• chemistry: The branch of science that deals with the composition and constitution of substances and the changes that they undergo as a consequence of alterations in the constitution of their molecules.

Key Terms

The Scientific Method

Classification of Matter

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• Matter can exist in one of three main states: solid, liquid, or gas.

• Solid matter is composed of tightly packed particles. A solid will retain its shape; the particles are not free to move around.

• Liquid matter is made of more loosely packed particles. It will take the shape of its container. Particles can move about within a liquid, but they are packed densely enough that volume is maintained.

• Gaseous matter is composed of particles packed so loosely that it has neither a defined shape nor a defined volume. A gas can be compressed.

Key Points

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• liquid: A substance that flows and keeps no definite shape because its molecules are loosely packed and constantly moving. It takes the shape of its container but maintains constant volume.

• gas: A substance that can only be contained if it is fully surrounded by a container (or held together by gravitational pull); a substance whose molecules have negligible intermolecular interactions and can move freely.

• solid: A substance that retains its size and shape without a container; a substance whose molecules cannot move freely except to vibrate.

Key Terms

Physical and Chemical Properties of Matter

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• All properties of matter are either physical or chemical, and physical properties are either intensive or extensive.

• Extensive properties, such as mass and volume, depend on the amount of matter measured.

• Intensive properties, such as density and color, do not depend on the amount of substance present.

• Physical properties can be measured without changing a substance’s chemical identity.

• Chemical properties can be measured only by changing a substance’s chemical identity.

​Key Points

Physical properties are properties that can be measured or observed without changing the chemical nature of the substance. Some examples of physical properties are:

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• color (intensive)

• density (intensive)

• volume (extensive)

• mass (extensive)

• boiling point (intensive): the temperature at which a substance boils

• melting point (intensive): the temperature at which a substance melts

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Remember, the definition of a chemical property is that measuring that property must lead to a change in the substance’s chemical structure. Here are several examples of chemical properties:

• Heat of combustion is the energy released when a compound undergoes complete combustion (burning) with oxygen. The symbol for the heat of combustion is ΔH

• Chemical stability refers to whether a compound will react with water or air (chemically stable substances will not react). Hydrolysis and oxidation are two such reactions and are both chemical changes.

• Flammability refers to whether a compound will burn when exposed to flame. Again, burning is a chemical reaction—commonly a high-temperature reaction in the presence of oxygen.

• The preferred oxidation state is the lowest-energy oxidation state that a metal will undergo reactions in order to achieve (if another element is present to accept or donate electrons).

Chemical Properties

• intensive property: Any characteristic of matter that does not depend on the amount of the substance present.

• extensive property: Any characteristic of matter that depends on the amount of matter being measured.

• physical property: Any characteristic that can be determined without changing the substance’s chemical identity.

• chemical property: Any characteristic that can be determined only by changing a substance’s molecular structure.

Key Terms

Measurement​s

​Standard Units (SI Units)

The International System of Units (abbreviated SI) is the metric system used in science, industry, and medicine.

​​Units of the SI System

There are seven base units in the SI system:

• the kilogram (kg), for mass

• the second (s), for time

• the kelvin (K), for temperature

• the ampere (A), for electric current

• the mole (mol), for the amount of a substance

• the candela (cd), for luminous intensity

• the meter (m), for distance

Derived units are based on units from the SI system of units.

​Here is a list of some commonly derived units:

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• Area: m²

• Volume: m³

• Velocity: m/s

• Acceleration: m/s²

• Density: g/mL or g/cm³

• Force: kg.m/s² or Newton (N)

• Energy: N.m or Joule (J)

Practice

Multiple Choice

Convert 300 meters to centimeters:

• 3000 cm

• 0.3 cm

• 30,000 cm

• 3 cn

Multiple Choice

1000 grams = _____ kilograms

• 0.0001
• 1,000,000
• 10
• 1

Multiple Choice

The metric unit for mass is 

• newton

• kilogram

• joule

• g/ml

Multiple Choice

Which lists metric units, in order, from smallest to largest?

• milligram, centigram, gram
• kilogram, gram, centigram
• decagram, hectogram, milligram
• kilogram, hectogram, decagram

Multiple Choice

7,000 g = ____ kg

• 70
• 700
• 7
• 0.07

Multiple Choice

The base unit for length is the

• second
• gram
• meter
• liter

Measurement Uncertainty​

Measurement uncertainty refers to the doubt that exists about the result of any measurement.

It quantifies the range within which the true value is expected to lie, taking into account potential errors or variations in the measurement process.

Uncertainty can arise from various sources, such as limitations in the measuring instrument, environmental conditions, or the skill of the person performing the measurement.

For example, if you measure the length of a table and report it as 100 cm ± 0.5 cm, the "± 0.5 cm" represents the measurement uncertainty, meaning the true length of the table could be anywhere between 99.5 cm and 100.5 cm.

Accuracy, Precision, and Error

Accuracy and precision are two important concepts in measurement that describe different aspects of measurement quality.

1. Accuracy refers to how close a measured value is to the true or accepted value. If a measurement is accurate, it means it is very close to the actual value. For example, if you know that a weight is exactly 50 grams, and you measure it as 49.9 grams, your measurement is accurate.

Accuracy and precision are two important concepts in measurement that describe different aspects of measurement quality.

1. Precision refers to how close repeated measurements are to each other, regardless of whether they are close to the true value. A precise measurement means that if you measure the same quantity multiple times, you get very similar results each time. For example, if you measure the weight of an object three times and get 49.9 grams, 50.0 grams, and 49.8 grams, your measurements are precise because they are close to each other, even if they are not exactly 50 grams.

To summarize:

• Accuracy: Closeness to the true value.

• Precision: Closeness of repeated measurements to each other.

A measurement can be accurate but not precise, precise but not accurate, both, or neither.

​Scientific Notation

Multiple Choice

Which of the following is correct scientific notation?

• 20.35 x 10⁴
• .2035 x 10⁴
• 2035 ⁴
• 2.035 x10⁴

Multiple Choice

Write 0.000972 in scientific notation.

• 9.72 x 10^-4

• 9.72 x 10^-3

• 972 x 10^-4

• 9.72 x 10⁴

Multiple Choice

Write 6,700 in scientific notation.

• 6.7 x 10³

• 6.7 x10^-3

• 67 x 10²

• 67 x 10^-2

Multiple Choice

Solve:
(6 x 10⁶) / (2 x 10³)

• 12 x 10³
• 3 x 10⁹
• 1.2 x 10⁴
• 3 x 10³

Multiple Choice

Which is the smallest?
1.3 x 10²⁰
2.9 x 10²¹
9.5 x 10³²
8.4 x 10¹⁹

• 1.3 x 10²⁰
• 2.9 x 10²¹
• 9.5 x 10³²
• 8.4 x 10¹⁹

Significant Figures

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​Significant figures are digits which contribute to the precision of a number.

​Key Points

• Significant figures are any non-zero digits or trapped zeros. They do not include leading or trailing zeros.

• When going between decimal and scientific notation, maintain the same number of significant figures.

• The final answer in a multiplication or division problem should contain the same number of significant figures as the original number with the fewest significant figures.

• In addition and subtraction, the final answer should contain the same number of decimal places as the original number with the fewest number of decimal places.

​Exact numbers are either defined numbers or the result of a count.

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For example, a dozen is defined as 12 objects, and a pound is defined as 16 ounces.

An exact number can only be expressed in one way and cannot be simplified any further.

​Exact numbers have an infinite number of significant figures, but they often appear as integers.

Exact Numbers

• Conversions within the American system (such as pounds to ounces, the number of feet in a mile, the number of inches in a foot, etc).

• Conversions with the metric system (such as kilograms to grams, the number of meters in a kilometer, the number of centimeters in a meter).

• The words per and percent means exactly out of one or one hundred, respectively.

• Counted numbers are exact: there are two chairs in the photograph. There are fifteen books on the shelf. Eighty-seven people attended the lecture.

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Examples of exact numbers include:

Measured numbers always have a limited number of significant digits.

A mass reported as 0.5 grams is implied to be known to the nearest tenth of a gram and not to the hundredth of a gram.

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There is a degree of uncertainty any time you measure something.

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For example, the weight of a particular sample is 0.825 g, but it may actually be 0.828 g or 0.821 g because there is inherent uncertainty involved.

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Measured Numbers

• The diameter of a coin, such as 10.2 mm.

• The weight of an object, such as 8.887 grams.

• The length of a pen, such as 12 cm.

Examples of measured numbers:

• systematic error: An inaccuracy caused by flaws in an instrument.

• Accuracy: The degree of closeness between measurements of a quantity and that quantity’s actual (true) value.

• Precision: Also called reproducibility or repeatability, it is the degree to which repeated measurements under unchanged conditions show the same results.

Key Terms

All measurements are subject to error, which contributes to the uncertainty of the result. Errors can be classified as human error or technical error.

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​Technical error can be broken down into two categories: random error and systematic error. Random error, as the name implies, occur periodically, with no recognizable pattern. Systematic error occurs when there is a problem with the instrument.

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For example, a scale could be improperly calibrated and read 0.5 g with nothing on it. All measurements would therefore be overestimated by 0.5 g.

Error