Scientific Method:

Observation: information obtained through the senses

Hypothesis: educated guess

Scientific Law: Rules

Variables in Experiments:

Dependent Variable: measurement

Independent Variable: what changes.

Controlled Variables: Kept constant during an experiment.

SI Units & Temperature Conversions:

Length: meter (m)

Mass: kilogram (kg)

Time: second (s)

Volume: Cubic meter (m^3)

Temp: Kelvin (K)

Amount of substance: mole

Electric current: ampere (A)

Accuracy vs. Precision:

Accuracy: How close measurement is to true value.

Precision: Consistency of measurements.

Percent error: estimated value and the actual value in comparison to the actual value and is expressed as a percentage.

Mixtures vs. Pure Substances:

1.Substances:

Substances have a definite, homogeneous composition, meaning they are uniform throughout.

Substances are further classified into:

• Elements: These are pure substances that consist of only one type of atom and are found on the periodic table. Examples include iron, oxygen, sulfur, and gold. Elements cannot be chemically separated into simpler substances.
• Compounds: These are pure substances composed of two or more elements that are chemically combined in fixed proportions. Examples include water (H₂O), sodium chloride (NaCl), and carbon dioxide (CO₂). Compounds can be chemically separated into their constituent elements.

2.Mixtures:

Mixtures have a variable composition and can be physically separated into their components.

 Mixtures are further divided into:

• Homogeneous mixtures (also called solutions): These have a uniform composition throughout, meaning you cannot distinguish the different components. Examples include air, steel, and salt water.
• Heterogeneous mixtures: These have distinct phases or parts, where the components are not uniformly distributed and can be visibly distinguished. Examples include soup and concrete.

Separation Techniques:

Decantation: Involves carefully pouring a liquid from a container, leaving any solid sediment behind. It is commonly used when the solid has settled at the bottom of the container.

Filtration: Separates solids from liquids or gases by passing the mixture through a filter, which allows the fluid to pass through while trapping the solid particles.

Simple Distillation: A process used to separate a liquid from a mixture by heating it until it vaporizes and then condensing the vapor back into a liquid. This method is ideal for separating substances with different boiling points.

Chromatography: A technique used to separate the components of a mixture based on their different affinities for a stationary phase (e.g., paper, gel) and a mobile phase (e.g., liquid, gas). The different components travel at different speeds, allowing for their separation.

Physical and Chemical Changes

Physical Changes:

A change that alters the substance without changing its composition. Often involves phase changes (e.g., boiling, freezing, melting).

Examples:

Crushing a can, melting ice, boiling water, shredding paper, and dissolving sugar.

Chemical Changes:

A process where one or more substances transform into new substances.Involves a chemical reaction.

Examples:

Rusting, burning, decomposing, and oxidizing.

Signs of Chemical Reactions:

1. Change in color.

2. Change in odor.

3. Formation of gas or vapor.

4. Energy change (heat or light).

5. Difficult to reverse.

Physical and Chemical Properties

Physical Properties:

Characteristics observed or measured without changing the substance.

 Types:

• Extensive: Depend on amount (e.g., mass, length, volume).
• Intensive: Independent of amount (e.g., density, color).

Examples:

Melting/boiling point, hardness, luster.

 Chemical Properties:

 The ability or inability of a substance to react with others.

 Examples:

 Iron rusts when exposed to oxygen.

Copper turns green in air.

Atomic theory

Dalton’s Atomic Model and Assumptions

1. Matter is made of atoms: All matter consists of tiny particles called atoms.

2. Atoms are indivisible: Atoms cannot be divided or destroyed, though later research showed they can be.

3. Atoms of different elements are different: Each element has its own type of atom, different from others.

4. Atoms of the same element are identical: All atoms of a specific element are the same in size, mass, and chemical properties.

5. Atoms combine in simple ratios: Atoms combine in fixed ratios to form compounds.

6. Atoms are rearranged in reactions: In chemical reactions, atoms are rearranged, but not created or destroyed.

Key Early Theories and Experiments

1. Democritus (circa 400 BCE)

• Theory: Atomism – Matter is made of indivisible particles called “atomos.”

• Importance: His idea set the stage for later atomic theories, though it was more philosophical than scientific.

2. John Dalton (1803)

• Theory: Atomic Theory – Matter is made of indivisible atoms that combine in simple ratios and are rearranged in chemical reactions.

• Importance: Dalton’s theory laid the foundation for understanding chemical reactions and the concept of atoms as fundamental particles.

3. J.J. Thomson (1897)

• Experiment: Discovery of the Electron – Using a cathode ray tube, Thomson discovered that atoms contain negatively charged particles called electrons.

• Model: Plum Pudding Model – Proposed that atoms are made of a positive “pudding” with negative electrons scattered inside, like plums.

4. Robert Millikan (1909)

• Experiment: Oil Drop Experiment – Measured the charge of the electron by observing oil droplets in an electric field.

• Importance: Millikan’s experiment helped determine the electron’s charge and confirmed its role in the atom.

5. Ernest Rutherford (1911)

• Experiment: Gold Foil Experiment – Discovered that atoms have a small, dense nucleus at the center, with electrons orbiting around it.

• Model: Nuclear Model – Proposed that most of the atom’s mass is concentrated in the nucleus, with the rest being empty space.

6. Niels Bohr (1913)

• Theory: Bohr’s Model of the Atom – Electrons orbit the nucleus in fixed energy levels, without radiating energy.

• Importance: Bohr’s model explained how electrons are arranged in atoms and how they gain or lose energy.

Periodic table

Atomic Number: Number of protons in an atom. It determines the element.

Mass Number: Total number of protons and neutrons in an atom.

Number of Protons: Equals the atomic number; positively charged particles in the nucleus.

Number of Electrons: Equals the number of protons in a neutral atom; negatively charged particles orbiting the nucleus.

Number of Neutrons: Mass number minus atomic number; neutral particles in the nucleus.

 Isotopes: atoms with the same number of protons but different numbers of neutrons or Atoms with the same atomic number but different mass numbers are called isotopes

Neutral Atom vs. Ion

Neutral Atom:

• Has equal numbers of protons and electrons.

• The positive charge of protons is balanced by the negative charge of electrons.

• Has no overall charge (net charge is 0).

Example:

A carbon atom (C) with 6 protons and 6 electrons is neutral.

Ion:

An atom that has gained or lost electrons, resulting in an overall net charge.

• Cations: Atoms that lose electrons, becoming positively charged (more protons than electrons).

• Anions: Atoms that gain electrons, becoming negatively charged (more electrons than protons).

Moles and Particles

1. Avogadro’s Number (Nₐ): The number of particles in one mole of a substance, 6.02 x 10^23 particles.

2. Mole (mol): A unit representing the amount of substance. One mole contains 6.02 x 10^23 particles.

3. Molar Mass (M): The mass of one mole of a substance, expressed in grams per mole (g/mol). For example, carbon (C) has a molar mass of 12.01 g/mol.

4. Representative Particles:

The smallest units of a substance:

• Atoms for elements

• Molecules for compounds

• Formula units for ionic compounds.