(CHE120) Chapter 1 Slides

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Chemistry

• Study of composition, properties, and interactions of matter

• Attempts to understand the behaviour of matter

• the “Central Science”

• Science based on observation & experimentation

Matter

Anything that takes up space & has mass.

What were the four elements believed by the ancient Greeks to constitute all matter?

Earth, air, fire, and water.

What was the primary goal of alchemists in early chemical history?

To transform base metals into noble metals.

What is one reason alchemy is not considered scientific by modern standards?

It lacked systematic experimentation and empirical validation.

Despite its limitations, what did alchemy contribute to the development of chemistry?

Useful ideas about manipulating matter.

This history of chemistry extends back over _____ years.

This history of chemistry extends back over 2500 years.

Provide 2 examples of how chemistry is used by humans in everday life.

• Digesting food

• Health and Medicine, e.g. vaccines, sanitation

• Synthesizing polymers for clothing, cookware, and credit cards

• Refining crude oil into gasoline and other products

• Energy and the environment, e.g. solar energy

Organic Chemistry

• Define

• Name 1 application

Def: The structure, properties, and reactions of carbon-containing compounds

Application: Drug synthesis

Inorganic Chemistry

• Define

• Name 1 application

Def: The structure, properties, and reactions of inorganic and organometallic compounds.

Application: Catalyst chemistry

Analytical Chemistry

• Define

• Name 1 application

Def: Quantitative measurements of composition and structure of matter.

Applications:

• Forensics

• Quality control

Biochemistry

• Define

• Name 1 application

Def: The chemistry of living systems

Examples of Subfields:

• Molecular genetics

• Protein science

Physical Chemistry

• Define

• Name 1 subfield

Def: Law of physics applied to chemical systems

Examples of Subfields:

• Kinetics

• Quantum chemistry

• Theoretical chemistry

What 2 things is chemistry fundamentally based on?

Observation & experimentation

Scientific Law

• Concise statement

• Summarizes several experimental observations & describes or predicts aspects of the natural world

Hypothesis

A tentative explanation for observations made during scientific research.

Theory v.s. Hypothesis

Theory

• Well-supported & tested explanation of a natural phenomenon

• Based on extensive evidence

Hypothesis

• Tentative explanation for observations made

• Testable, falsifiable statement

• Predicts an outcome

Theory v.s. Scientific Law

Law = Statement about observed phenomenon

• E.g: Law of Conservation of Mass summarizes the observation that mass can not be destroyed nor created.

Theory = Explanation of hypothesis that’s been tested & well-supported

• E.g: Dalton’s Atomic Theory explains the nature of atoms and their behavior.

Substance

Form of matter that has a definite composition & distinct properties.

Mass v.s. Weight

Mass = Measure of amount of matter (particles) in an object

Weight = Force exerted by gravity on an object

For example:

• on earth → your mass is 150lbs, your weight is 150lbs

• on moon → your mass is 150lbs, but your weight will change as gravity changes

T or F: The mass of an object/being changes depending on their location in the universe.

False.

• The mass of all objects/beings remains constant regardless of their location

• What does change w/ location is weight → bc weight depends on the gravitational force of the celestial body the object/being is on

The mass of an object/being only changes if matter is _______ or _______.

The mass of an object/being only changes if matter is added or removed.

Astronaut Paul weighs 180 lbs on the Moon. What will happen to his weight when he comes back to Earth?

His weight will increase since Earth has higher gravity than the Moon.

The 3 Phases of Water

• Phase of Matter

• Arrangement of Molecules

• Example

1) Solid (ice)

• Molecules → Close together & organized

• E.g. Icebergs

2) Liquid (liquid water)

• Molecules → Close together & disorganized

• E.g. Lakes, oceans

3) Gas (water vapour)

• Molecules → Far apart & very disorganized

• E.g. Moisture in air

What do the symbols (g), (s), and (l) represent?

(g) → gas

(s) → solid

(l) → liquid

Describe the 3 phases of matter.

(1) Solids

• Definite shape

• Definite volume

• Not compressible

(2) Liquids

• No definite shape

• Definite volume

• Able to flow

(3) Gases

• No definite shape

• No definite volume

• Molecules are very far apart

• Highly compressible

__________ take the shape of the container they’re in while maintaining a constant volume.

Liquids take the shape of the container they’re in while maintaining a constant volume.

In molecular models, what do red spheres represent?

Oxygen atoms

In molecular models, what do white spheres represent?

Hydrogen atoms

Briefly explain how substances can be changed from one phase of matter to another.

Substances can be changed from one phase of matter to another by adding or removing energy — usually heat.

Adding energy

• particles move faster

• substance will eventually melt or evaporate

  • melt = solid → liquid

  • evaporate = liquid → gas

Removing energy

• particles slow down

• substance will eventually freeze or condensate

  • freeze = liquid → solid

  • condensate = gas → liquid

Describe the Classifications of Matter

The Two Major Classifications of Matter

(1) Mixture = 2 or more substances that CAN be physically separated

(2) Pure Substance = 2 or more substances that CAN’T be physically separated

Two Types of Mixtures

i) Homogeneous

• AKA a solution

• Uniform composition throughout (which is also visible)

• E.g. sprite

ii) Heterogenous

• Non-uniform composition

• The different phases of matter are visible

• E.g. vinaigrette dressing (oil + non-oil substance; they are visibly separate)

Two Types of Pure Substances

i) Compounds

• Pure substance composed of 2+ elements that CAN be chemically decomposed

• E.g. Mercury (II) Oxide decomposes into mercury & oxygen when heated

  

ii) Element

• Pure substance that CANNOT be chemically physically decomposed

• Examples:

  • Oxygen (O)

  • Gold (Au)

  • Helium (He)

What is the smallest particle that retains the properties of an element?

An atom

T or F: Atoms can be chemically broken down further while retaining their elemental identity.

False.

An ____ is the most basic unit of chemical matter.

An atom is the most basic unit of chemical matter.

What are atoms composed of?

• Protons

• Neutrons

• Electrons

T or F: Molecules can be separated into individual atoms.

True. Through chemical reactions

Name 2 examples of molecules made from a single element.

• O₂ (oxygen gas)

• N₂ (nitrogen gas)

Atoms v.s. Molecules

Atoms (of an element)

• Most basic unit of chemical matter

• Smallest type of particle that can have properties of an element

• Cannot be broken down further while retaining properties of the element

Molecules (of an element)

• Combination of 2 or more atoms

• Can be broken down into individual atoms

Elements v.s. Compounds

Elements & compounds are the two major categories of pure substances.

Elements

• Simplest type of pure substance

• Consists of only 1 type of atom

• Can’t be broken down further by physical or chemical means

• Forms :

  • Single atoms

    • e.g. Gold (Au); Helium (He); Iron (Fe)

  • Molecules of an element — still considered elements because all atoms are of the same type.

    • e.g., O₂, N₂

Compounds

• Pure substances composed of 2+ types of atoms that are chemically bonded

• Can be broken down chemically into elements

• Compounds have diff properties than the elements that form it

• Examples: Water (H₂O), Carbon dioxide (CO₂), Sodium chloride (NaCl).

In one sentence, state the critical difference b/w elements & compounds.

Elements → Cannot be broken down chemically

Compounds → Can be broken down chemically into elements

Compound

Pure substance composed of 2+ elements

How many elements in the periodic table?

118 elements

Each element on the periodic table is identified by its _______________. The atomic number represents the _______________ in the element’s nucleus.

Each element on the periodic table is identified by its atomic number. The atomic number represents the number of protons in the element’s nucleus.

Natural Elements

• Define

• 2 Examples

Def: Elements that naturally occur on Earth

(Elements number 1 through 94 on periodic table)

Examples:

• Gold (Au)

• Oxygen (O)

• Carbon (C)

• Sulfur (S)

Synthetic Elements

• Define

• How can they be identified on periodic table?

• 2 Examples

Def: Man-made in labs through nuclear reactions

Identify by their atomic number → elements with number 95 through 118 on periodic table

Examples:

• Americium (Am)

• Seaborgium (Sg)

Discuss the historical context of the periodic table and how its evolved over centuries.

• Ancient elements: Known since antiquity (e.g., gold, copper).

• Lavoisier’s list (1789): Early classification of known elements.

• Mendeleev’s table (1869): Organized elements by atomic mass and predicted undiscovered ones.

• Seaborg’s table (1945): Included actinides and expanded the table.

• Modern updates: Continued discoveries up to 2012 and beyond.

How many elements naturally occur on Earth?

94

How many elements are man-made (synthetic)?

24

________________ are elements that are created through nuclear reactions.

Synthetic elements are elements that are created through nuclear reactions.

Give 2 examples of natural elements.

• Iron (Fe)

• Oxygen (O)

• Nitrogen (N)

• Phosphorus (P)

Give 2 examples of synthetic elements.

Americium (Am) and Seaborgium (Sg).

In the periodic table the first ___ elements are _________, whereas the rest are _________.

In the periodic table the first 94 elements are natural, whereas the rest are synthetic.

Who published the first organized periodic table in 1869?

Dmitri Mendeleev.

What defines an element’s position on the periodic table?

It’s atomic number

What was Seaborg’s major contribution to the periodic table?

Inclusion of the actinide series and expansion of synthetic elements.

When the compound Mercury II Oxide is __________ it decomposes into two elements: ___________ and __________.

When the compound Mercury II Oxide is heated it decomposes into two elements: liquid mercury and oxygen gas.

Physical Properties v.s. Chemical Properties

• Compare definitions

• State the 5 most common examples/scenarios for each

Physical Property

Can be observed w/o changing composition of substance

Most common examples:

1. State/Phase of Matter — solid v.s. liquid v.s. gas at room temp

2. Boiling/Melting Point — temp at which phase change occurs

3. Density — Relationship b/w mass & volume (D = m / V)

4. Malleability — Ability to be re-shaped w/o breaking

5. Solubility — How well a substance dissolves in a solvent

Chemical Property

Describes how likely a substance is to undergo a chemical change that transforms it into a new/different substance,

Most common examples:

1. Flammability — How easily substance ignites/burns

2. Reactivity — How a substance reacts with water, acids, or oxygen (e.g. rusting)

3. Toxicity — Potential for substance to damage an organism

4. Acidity/Alkalinity (pH) — How a substance behaves as an acid or a base

Chemical Property v.s. Chemical Change.

Provide an example for each.

Chemical Property

• A characteristic or "personality trait" of a substance.

• Think “Ability”

• E.g. Wood is flammable

Chemical Change

• An event where a substance actually transforms.

• Think “Process” or “Event”

• E.g. Wood is burning

Physical Change v.s. Chemical Change

• Compare definitions

• State the 5 most common examples/scenarios for each

Physical Change

• Change in substance’s state/phase

• E.g. Ice melting

Chemical Change

• Chemical reaction

• E.g. rusting, burning, baking

Condensation

High-Yield

DEF: Phase change of a water vapour into a liquid

Not from slides (extra info to read for better understanding)

• Occurs when warm, moist air cools to below its dew point

• As water vapour molecules cool, their kinetic energy decreases

• Results in their motion slowing down enough for intermolecular attractions to pull them together close enough to form liquid water droplets on a surface or in the air

Evaporation

High-Yield

DEF: Phase change of a liquid water into water vapour (gas)

Not from slides (extra info to read for better understanding)

• Phase change of liquid water into water vapor (gas).

• Occurs when surface molecules of a liquid gain enough energy to escape into the air.

• Happens below the boiling point, often when air is warm, dry, or moving.

• Molecules absorb energy, their kinetic motion increases, and they overcome intermolecular forces.

• High‑energy molecules leave first, causing cooling of the remaining liquid.

Name & briefly describe the components of a measurement.

Every measurement includes:

1. Number (size/magnitude)

2. Unit (standard of comparison)

3. Uncertainty (estimated last digit)

Example: 4.7 ± 0.1 cm

• Number: 4.7

• Unit: cm

• Uncertainty: 0.1

Without units, the number in a measurement is ________.

Without units, the number in a measurement is meaningless.

What are SI units? Explain its importance in the world of science.

• SI = International System of Units

• Most widely used measurement system worldwide

• Adopted in 1964

• Modernized, globally agreed-upon version of metric system

Compare the 2 systems of measurement used in the USA.

U.S. Customary System (USCS)

Used for:

• Daily life

• Commerce

Includes:

• Length → inches, feet, yards, miles

• Weight → ounces, pounds, tons

• Volume → Gallons, quarts, pints, cups

Internal System of Units

Used for:

• Science

• Medicine

• Internal trade

For each of the 7 base SI units, provide:

• Property

• Name of Unit

• Symbol of Unit

• What It Represents

• Why It's a Base Unit

Property	Name of Unit	Symbol of Unit	What It Represents	Original Definition	Current Definition
Length	Meter	m	Distance b/w 2 points	1/10,000,000 of distance from North Pole to equator	Distance light travels in a vacuum in 1/299,792,458
Mass	Kilogram	kg	Amount of matter	Mass of a litre of water	Defined in terms of the Planck constant, h, which states
Time	Second	s	Duration of events
Temperature	Kelvin	K	Average kinetic energy of particles
Electric Current	Ampere	A	Flow of electric charge
Amount of Substance	Mole	mol	Number of particles
Luminous Intensity	Candela	cd	Received brightness

Each of the 7 base units were selected because they correspond to …

to a fundamental physical property that cannot be broken down into something more basic.

Which 2 other countries besides the USA have not officially adopted SI units?

1. Liberia

2. Myanmar

What is the SI base unit for length?

Meter (m)

What is the SI base unit for mass?

Kilogram (kg)

What is the SI base unit for time?

Second (s)

What is the SI base unit for temperature?

Kelvin (K)

What is the SI base unit for electric current?

Ampere (A)

What is the SI base unit for amount of substance?

Mole (mol)

What is the SI base unit for luminous intensity?

Candela (cd)

What is Kelvin (K)?

• Absolute temperature scale

• Has no degree symbol

What is Celsius (°C)?

Common scientific temperature scale

What is the formula to convert °C to °F?

What is the formula to convert °C to K?

SI unit for volume

Cubic meter (m³)

What is 1 dm³ equal to?

1 L

What is 1 cm³ equal to?

1 mL

Extensive Properties vs. Intensive Properties

Extensive Properties

• Dependent on amount of matter

• Changes when sample size changes

• Examples:

  • Mass

  • Volume

  • Total energy

Intensive Properties

• Independent of amount of matter

• Stays constant regardless of sample size

• Examples:

  • density

  • colour

  • melting point

  • temperature

T or F: The periodic table organizes elements into groups with similar properties.

True

Elements in the same group (vertical column) share similar chemical properties (bc they all have same number of valence electrons).

What is the formula density?

Density = mass / volume

What are common density units for solids and liquids?

g/cm³

What are common density units for gases?

g/L

How is scientific notation written?

N × 10ⁿ

What must N be between in scientific notation?

What must N be between in scientific notation?

Why is scientific notation used?

For very large or very small numbers.

What is Avogadro’s number in scientific notation?

6.022 × 10²³

What does kilo (k) represent?

10³

What does centi (c) represent?

1A: 10⁻²

0⁻²

What does milli (m) represent?

10⁻³

What is accuracy?

Closeness to the true value.

What is precision?

Reproducibility of measurements.

What are exact numbers?

Counting or defined quantities; infinite significant figures.

What are uncertain numbers?

Measured values that include estimation.

What digits are included in significant figures?

All certain digits plus one uncertain digit.

Which digits are always significant?

Nonzero digits.