Physical science 11

Chemistry as a Science

• Definition: Chemistry is a physical science that studies the properties and interactions of matter.

Historical Context

• Ancient Contributions:

  • Civilizations:

  • Egyptian, Chinese, and Mesopotamian civilizations practiced early forms of chemistry such as wine making, metallurgy, dyeing, glass making, pottery, and embalming fluids.

  • Date: Earliest records date back to 3500 BC.

• Alchemy (500 – 1600 AD):

  • Main Objectives:

  • Transformation of common metals into gold.

  • Discovery of an elixir of life.

• Modern Chemistry:

  • Beginning: Established in 1774 by French chemist Antoine Lavoisier.

  • Methodology: Emphasized the use of quantitative methods while avoiding mysticism, superstition, and secrecy.

Major Divisions of Chemistry

1. Physical Chemistry:

   • Applies theories of physics to chemical systems.

2. Analytical Chemistry:

   • Identifies the composition (what substances are present) and quantity (how much) of substances.

3. Organic Chemistry:

   • Focuses on carbon-containing compounds.

4. Inorganic Chemistry:

   • Concerns non-carbon compounds.

5. Biochemistry:

   • Studies chemical reactions that occur in living organisms.

Classification of Matter

• Matter: Anything that has mass and occupies space.

  • Types:

  1. Pure Substances:

     • Composition remains constant across samples.

     • Can be further divided into:

       • Elements:

       • Substances composed of only one type of atom. Cannot be broken down further chemically. There are 114 known elements.

       • Compounds:

       • Pure substances made of two or more elements chemically combined in fixed proportions.

  2. Mixtures:

     • Compositions can vary.

     • Can be further classified into:

       • Homogeneous Mixtures: Similar compositions (e.g., solutions).

       • Heterogeneous Mixtures: Dissimilar compositions (e.g., concrete).

Chemical Elements and Atoms

• Elements:

  • Defined as substances that cannot be broken down further through chemical reactions.

• Atom:

  • The smallest particle of an element that retains the element's chemical properties.

Molecules

• Molecule:

  • The smallest particle of a compound that retains its chemical properties.

  • Compounds have distinct properties that differ from their constituent elements.

  • Example: Water ($H<em>2O$) can be broken down into $H</em>2$ and $O_2$ through electrolysis.

Classification of Mixtures

• Mixtures: Substances that consist of two or more components physically mixed rather than chemically bonded.

  • Separation: Mixtures can be separated through physical means (e.g. heating).

  • Examples: Ice, Pepto Bismol, Oxygen gas in air, Alcoholic drinks, Copper wiring, Sugar.

Solutions

1. Liquid Solutions:

   • Solvent: The major component of a solution, usually in greater quantity.

   • Solute: The substance that is dissolved in the solvent.

   • Aqueous Solution: A solution where water is the solvent.

2. Types of Solutions:

   • Unsaturated Solution: Can dissolve more solute at the current temperature.

   • Saturated Solution: Contains the maximum amount of solute that can be dissolved in the solvent at the current temperature.

   • Supersaturated Solution: Contains more solute than is typically soluble at that temperature.

Solubility of Gases

• Solubility increases with pressure (e.g. CO2 in soft drinks) and decreases with temperature.

• Example: CO2 is forced into beverages under pressure and escapes when opened due to atmospheric pressure reduction.

Dalton's Atomic Theory

1. All matter consists of tiny particles called atoms.

2. All atoms in an element have identical chemical properties.

3. Atoms combine in whole number ratios to form compounds.

4. Atoms cannot be created or destroyed in chemical reactions (law of conservation of matter).

• Significance: Dalton's atomic theory has persisted for over 200 years and has been supported by advancements in atomic observation.

Structure of Atoms

• Components of Atoms:

  • Protons: Positively charged particles that define the type of atom (atomic number).

  • Neutrons: Neutrally charged particles that can vary within an atom.

  • Electrons: Negatively charged particles that orbit the nucleus.

• Properties:

  • The nucleus contains 99% of an atom's mass but occupies only 1% of its volume.

  • Atomic Number (Z): Number of protons; unique to each element.

  • Mass Number (A): Total number of protons and neutrons.

  • Charge: Resulting from the balance of protons minus electrons.

Isotopes and Atomic Mass

• Isotope: Atoms with the same atomic number but different neutron counts.

• Natural Abundance: Percentage occurrence of each isotope.

• Mass Spectrometry: A technique used for weighing and identifying isotopes.

• Units of Measure: Atomic mass units (amu) can be converted to grams.

• Avogadro’s Number: $N_{A}=6.022times10^{23}{ atoms/mole}$  allowing for counting moles of atoms based on molar mass.

Periodic Table Organization

• Elements are arranged by increasing atomic number and share similar chemical properties.

• Categories:

  • Metals: Good conductors, malleable, ductile, typically solid at room temperature.

  • Nonmetals: Brittle, poor conductors, many gaseous at room temperature.

  • Metalloids: Exhibit properties of both metals and nonmetals.

• Geographical Layout:

  • Groups: metals on the right and nonmetals towards the left 

Chemical Formulas and Compounds

• Chemical Formulas: Used to represent compounds containing various elements.

• Monoatomic and Diatomic Elements:

  • Monoatomic: Elements existing as single atoms.

  • Diatomic: Elements that form pairs

  • Polyatomic: Elements with multiple bonded atoms 

Naming Conventions for Compounds

1. Binary Compounds:

   • Only two elements except hydrogen. 

   • Example: $CO_2$ is carbon dioxide, where prefixes denote the number of atoms.

   • lower proton number gets written first 

2. Ionic Compounds:

   • Formed by cations and anions, named similar to nonmetallic compounds without prefixes.

   • Include polyatomic ion notations in parentheses if needed.

Allotropes 

Allotrope – two or more forms of the same element that have different bonding structures in the same physical phase
Example: Diamond, Graphite (and fullerenes)
- Both pure Diamond and pure Graphite are each 100% carbon (C),
and are both solid
- But the atomic arrangement of the carbon atoms is different

Reactions and Ions

• Ionic Compounds: Formed when cations (positives) and anions (negatives) bond.

• Common Ions:

  • Group 1: +1, Group 2: +2, Group 17: -1, etc.

• Ions and Charge Conservation:

  • Ions are formed through the loss or gain of electrons creating cations (+ charged) or anions (- charged).

Ionization Energy

• Definition: The energy required to remove an electron from an atom increases across a period and decreases down a group due to the electrons’ proximity to the nucleus.

Acid Nomenclature

• Acids: Formed when hydrogen combines with polyatomic ions.

  • Examples of Acids:

  • Hydrochloric acid ($HCl(aq)$)

  • Nitric acid ($HNO_3(aq)$)

  • Sulfuric acid ($H<em>2SO</em>4(aq)$)

  • Phosphoric acid ($H<em>3PO</em>4(aq)$)

  • Acetic acid ($HC<em>2H</em>3O_2(aq)$)

Overall, this comprehensive overview of chemistry includes key historical insights, foundational definitions, and important concepts such as atomic theory, chemical bonding, nomenclature, and the periodic table, acting as an essential guide for understanding the subject.

Electron Configuration

Electrons are located in energy levels or shells that surround
the nucleus
 Level 1 – maximum of 2 electrons
 Level 2 – maximum of 8 electrons
 Level 3 – maximum of 18 electrons
 The chemical reactivity of the elements depends on the order of
electrons in these energy levels
 The outer shell of an atom is known as the valence shell
 The electrons in the outer shell are called the valence electrons
 The valence electrons are the electrons involved in forming
chemical bonds – so they are extremely important
 Elements in a given group all have the same number of valence
electrons (and similar chemical properties)

The number of electrons in an atom is the same as the element’s
atomic number (Z)
 The number of shells that contain electrons will be the same as
the period number that it is in

Electron Shell Distribution

• Atoms are neutral when the number of protons and electrons are
  equal.

• Atoms or molecules with a charge are called ions.

• If you lose an electron, there is now more protons and so the atom
  now has a positive charge (cation)

• If an atom gains an electron, it now has more electrons than protons
  and so the atom will have an overall negative charge (anion)

Like charges repel each other and opposite charges attract each other, so cations
attract anions to make a collection of cations and anions

Solids that contain cations and anions in balanced whole number ratios are called
ionic compounds (salts

Know these

Group 1 - +1
Group 2 - +2
Group 13 - +3
Group 14 - +4 to -4
Group 15 - -3
Group 16 - -2
Group 17 - -1
Group 18 - 0

The Periodic Nature of Atomic Size

The atomic size of the elements also varies periodically
 Atomic size increases down a group
 Atomic size decreases across a period
 The atoms on the far left are the largest due to less charge
(fewer protons) in the nucleus and  the outer electrons are
more loosely bound

Ionization Energy-Also periodic

Ionization energy – the amount of energy that it takes to remove
an electron from an atom
 Ionization energy increases across a period due to additional
protons in the nucleus
 Ionization energy decreases down a group because of the
additional shells situated between the nucleus and the outer
electron shell.

Remember all the polyatomic ions

• 1) Ionic compounds are neutral overall so the cationic and anionic charges
  need to balance out.
  Other rules:
  2) the cation is always listed before the anion.
  3) the formula of any polyatomic ion is written as a unit.
  4) polyatomic ions are placed in parentheses with a subscript to
  indicate ratios different from 1:1.
  
  

Exam Topics

Classification of Matter

• Matter: Anything that has mass and occupies space.

  • Types:

    1. Pure Substances:

       • Composition remains constant across samples.

       • Can be further divided into:

         • Elements:

           • Substances composed of only one type of atom. Cannot be broken down further chemically. There are 114 known elements.

         • Compounds:

           • Pure substances made of two or more elements chemically combined in fixed proportions.

    2. Mixtures:

       • Compositions can vary.

       • Can be further classified into:

         • Homogeneous Mixtures: Similar compositions (e.g., solutions).

         • Heterogeneous Mixtures: Dissimilar compositions (e.g., concrete).

Chemical Elements and Atoms

• Elements:

  • Defined as substances that cannot be broken down further through chemical reactions.

• Atom:

  • The smallest particle of an element that retains the element's chemical properties.

Molecules

• Molecule:

  • The smallest particle of a compound that retains its chemical properties.

  • Compounds have distinct properties that differ from their constituent elements.

  • Example: Water ($H<em>2O$) can be broken down into $H</em>2$ and $O_2$ through electrolysis.

Solutions

1. Liquid Solutions:

   • Solvent: The major component of a solution, usually in greater quantity.

   • Solute: The substance that is dissolved in the solvent.

   • Aqueous Solution: A solution where water is the solvent.

2. Types of Solutions:

   • Unsaturated Solution: Can dissolve more solute at the current temperature.

   • Saturated Solution: Contains the maximum amount of solute that can be dissolved in the solvent at the current temperature.

   • Supersaturated Solution: Contains more solute than is typically soluble at that temperature.

Solubility of Gases

• Solubility increases with pressure (e.g. CO2 in soft drinks) and decreases with temperature.

• Example: CO2 is forced into beverages under pressure and escapes when opened due to atmospheric pressure reduction.

Chemical Formulas and Compounds

• Chemical Formulas: Used to represent compounds containing various elements.

• Monoatomic and Diatomic Elements:

  • Monoatomic: Elements existing as single atoms.

  • Diatomic: Elements that form pairs

  • Polyatomic: Elements with multiple bonded atoms

Periodic Table Organization

• Elements are arranged by increasing atomic number and share similar chemical properties.

• Categories:

  • Metals: Good conductors, malleable, ductile, typically solid at room temperature.

  • Nonmetals: Brittle, poor conductors, many gaseous at room temperature.

  • Metalloids: Exhibit properties of both metals and nonmetals.

• Geographical Layout:

  • Groups: metals on the right and nonmetals towards the left

Electron Configuration

Electrons are located in energy levels or shells that surround the nucleus.

• Level 1 – maximum of 2 electrons

• Level 2 – maximum of 8 electrons

• Level 3 – maximum of 18 electrons

• The chemical reactivity of the elements depends on the order of electrons in these energy levels.

• The outer shell of an atom is known as the valence shell.

• The electrons in the outer shell are called the valence electrons.

• The valence electrons are the electrons involved in forming chemical bonds – so they are extremely important.

• Elements in a given group all have the same number of valence electrons (and similar chemical properties).
  The number of electrons in an atom is the same as the element’s atomic number (Z).

• The number of shells that contain electrons will be the same as the period number that it is in.

Electron Shell Distribution

• Atoms are neutral when the number of protons and electrons are equal.

• Atoms or molecules with a charge are called ions.

• If you lose an electron, there is now more protons and so the atom now has a positive charge (cation).

• If an atom gains an electron, it now has more electrons than protons and so the atom will have an overall negative charge (anion).
  Like charges repel each other and opposite charges attract each other, so cations attract anions to make a collection of cations and anions.
  Solids that contain cations and anions in balanced whole number ratios are called ionic compounds (salts).
  Know these
  Group 1 - +1
  Group 2 - +2
  Group 13 - +3
  Group 14 - +4 to -4
  Group 15 - -3
  Group 16 - -2
  Group 17 - -1
  Group 18 - 0

The Periodic Nature of Atomic Size

• The atomic size of the elements also varies periodically.

• Atomic size increases down a group.

• Atomic size decreases across a period.

• The atoms on the far left are the largest due to less charge (fewer protons) in the nucleus and $\therefore$ the outer electrons are more loosely bound.

Ionization Energy-Also periodic

• Ionization energy – the amount of energy that it takes to remove an electron from an atom.

• Ionization energy increases across a period due to additional protons in the nucleus.

• Ionization energy decreases down a group because of the additional shells situated between the nucleus and the outer electron shell.

Naming Conventions for Compounds

1. Binary Compounds:

   • Only two elements except hydrogen.

   • Example: $CO_2$ is carbon dioxide, where prefixes denote the number of atoms.

   • lower proton number gets written first

2. Ionic Compounds:

   • Formed by cations and anions, named similar to nonmetallic compounds without prefixes.

   • Include polyatomic ion notations in parentheses if needed.

Acid Nomenclature

• Acids: Formed when hydrogen combines with polyatomic ions.

  • Examples of Acids:

    • Hydrochloric acid ($HCl(aq)$)

    • Nitric acid ($HNO_3(aq)$)

    • Sulfuric acid ($H<em>2SO</em>4(aq)$)

    • Phosphoric acid ($H<em>3PO</em>4(aq)$)

    • Acetic acid ($HC<em>2H</em>3O_2(aq)$)

Remember all the polyatomic ions

1) Ionic compounds are neutral overall so the cationic and anionic charges need to balance out.
Other rules:
2) the cation is always listed before the anion.
3) the formula of any polyatomic ion is written as a unit.
4) polyatomic ions are placed in parentheses with a subscript to indicate ratios different from 1:1.