Y10 Structure and Bonding 1 (Edexcel) PDF

**Structure and Bonding 1**\ **(Topic No. 3.1)** **Formation of Ionic Bonds** - **Ionic Bonds**: A type of chemical bond formed through the transfer of electrons from one atom to another. - **Atoms**: The basic units of matter, made up of protons, neutrons, and electrons. - **Elect...

**Structure and Bonding 1**\ **(Topic No. 3.1)** **Formation of Ionic Bonds** - **Ionic Bonds**: A type of chemical bond formed through the transfer of electrons from one atom to another. - **Atoms**: The basic units of matter, made up of protons, neutrons, and electrons. - **Electrons**: Negatively charged particles that orbit the nucleus of an atom. - **Cations**: Positively charged ions formed when an atom **loses electrons**. - Example: Sodium (Na) loses one electron to become Na⁺. - **Anions**: Negatively charged ions formed when an atom **gains electrons**. - Example: Chlorine (Cl) gains one electron to become Cl⁻. - **Formation of Ionic Bonds**: - **Electron Transfer**: An electron moves from a metal atom (which becomes a cation) to a non-metal atom (which becomes an anion). - This transfer creates opposite charges, resulting in an **electrostatic attraction** between the cation and anion. - **Dot and Cross Diagrams**: A visual representation used to show the transfer of electrons and the formation of ionic bonds. - **Dots** represent the **valence electrons** of one atom. - **Crosses** represent the **valence electrons** of another atom. - When drawing: - Start with the electron configurations of both atoms. - Show the transfer of electrons by moving a dot to the cross side or vice versa. - Indicate the formation of cations and anions clearly. - **Example of Sodium Chloride (NaCl)**: - **Sodium (Na)** has 1 electron in its outer shell (1 electron → cation Na⁺). - **Chlorine (Cl)** has 7 electrons in its outer shell (gains 1 electron → anion Cl⁻). - **Ionic bond** forms between Na⁺ and Cl⁻, resulting in NaCl (table salt). Fig.: Formation of Sodium Chloride (NaCl) - **Properties of Ionic Compounds**: - High melting and boiling points. - Conduct electricity when dissolved in water or melted. - Form crystalline structures. **Ions** - **Ion Definition**: - An **ion** is an atom or a group of atoms that has a **positive** or **negative charge**. - **Charge Explanation**: - **Positive charge**: When an ion has more **protons** than **electrons**, it is called a **cation**. - **Negative charge**: When an ion has more **electrons** than **protons**, it is called an **anion**. - **Formation of Ions**: - Ions are formed when atoms **gain** or **lose** electrons. - **Gaining electrons** leads to a negative charge (anion). - **Losing electrons** leads to a positive charge (cation). - **Examples of Ions**: - Common **cations**: - Sodium ion (**Na⁺**) - Calcium ion (**Ca²⁺**) - Common **anions**: - Chloride ion (**Cl⁻**) - Sulfate ion (**SO₄²⁻**) - **Importance of Ions**: - Ions play a crucial role in **chemical reactions** and **biological processes**. - They are essential for the **formation of compounds** and can affect the **properties** of substances. - **Ionic Compounds**: - When cations and anions combine, they form **ionic compounds**. - Example: Sodium chloride (**NaCl**) is formed from sodium ions (**Na⁺**) and chloride ions (**Cl⁻**). - **Conductivity**: - Ions are important in **conductivity**: - Ions in solution can carry **electric current**, making them vital in applications like batteries and electrolysis. - **Electrolytes**: - Substances that dissociate into ions in solution are known as **electrolytes**. - Common electrolytes include **salts**, **acids**, and **bases**. - They are essential for various bodily functions, including **nerve transmission** and **muscle contraction**. - **Charge Balance**: - In a neutral compound, the total positive charge of cations equals the total negative charge of anions, maintaining **charge balance**. **Calculations involving simple ions** - **Atoms** are made up of three types of subatomic particles: **protons**, **neutrons**, and **electrons**. - The **atomic number** of an element indicates the number of **protons** in its nucleus. - The **mass number** is the total number of protons and **neutrons** in the nucleus of an atom. - To find the number of **neutrons** in an atom: - Use the formula: **Mass Number - Atomic Number = Number of Neutrons** - **Electrons** are usually equal to the number of protons in a **neutral atom**. However, in **ions**, this changes: - **Cations**: Positively charged ions, have lost electrons. - Example: Na⁺ (Sodium ion) - Atomic Number = 11 (11 protons) - Mass Number = 23 - Neutrons: 23 - 11 = 12 - Electrons: 11 - 1 (lost 1 electron) = 10 - Summary: 11 protons, 12 neutrons, 10 electrons. - **Anions**: Negatively charged ions, have gained electrons. - Example: Cl⁻ (Chloride ion) - Atomic Number = 17 (17 protons) - Mass Number = 35 - Neutrons: 35 - 17 = 18 - Electrons: 17 + 1 (gained 1 electron) = 18 - Summary: 17 protons, 18 neutrons, 18 electrons. - **Steps to Calculate for Simple Ions**: - Identify the **atomic number** (Z) and **mass number** (A). - Determine the number of **protons** (equal to atomic number). - Calculate the number of **neutrons** using: A - Z. - Determine the number of **electrons** based on the ion\'s charge: - For cations, subtract the charge from the number of protons. - For anions, add the charge to the number of protons. **Formation of Ionic Compounds** - **Ionic Compounds**: Formed from the interaction between **cations** (positively charged ions) and **anions** (negatively charged ions). - **Formation of Ions**: - Atoms can lose or gain **electrons** to achieve a full outer shell, similar to the noble gases. - **Groups 1 and 2** (alkali and alkaline earth metals): - **Group 1 (Alkali Metals)**: - Have **1 electron** in their outer shell. - Tend to **lose 1 electron** to form **+1 cations** (e.g., **Na⁺**, **K⁺**). - **Group 2 (Alkaline Earth Metals)**: - Have **2 electrons** in their outer shell. - Tend to **lose 2 electrons** to form **+2 cations** (e.g., **Ca²⁺**, **Mg²⁺**). - **Groups 6 and 7** (chalcogens and halogens): - **Group 6 (Chalcogens)**: - Have **6 electrons** in their outer shell. - Tend to **gain 2 electrons** to achieve a full outer shell, forming **-2 anions** (e.g., **O²⁻**, **S²⁻**). - **Group 7 (Halogens)**: - Have **7 electrons** in their outer shell. - Tend to **gain 1 electron** to form **-1 anions** (e.g., **Cl⁻**, **F⁻**). - **Electron Transfer**: - Ionic compounds form when **electrons** are transferred from metals (groups 1 and 2) to nonmetals (groups 6 and 7). - This transfer creates charged ions that attract each other due to **electrostatic forces**. - **Examples**: - **Sodium Chloride (NaCl)**: - **Na** (group 1) loses 1 electron to become **Na⁺**. - **Cl** (group 7) gains 1 electron to become **Cl⁻**. - They combine to form NaCl. - **Calcium Oxide (CaO)**: - **Ca** (group 2) loses 2 electrons to become **Ca²⁺**. - **O** (group 6) gains 2 electrons to become **O²⁻**. - They combine to form CaO. - **Endings in Compound Names**: - Compounds often have specific endings that indicate their chemical composition. - **Ending --ide**: - Used for **binary compounds** (compounds comprised of two elements). - Typically indicates that the compound contains a **nonmetal**. - Examples: - **NaCl**: Sodium Chloride (table salt) - **H~2~O**: Water (Dihydrogen Oxide) - Generally, the name is formed by taking the name of the first element and adding "-ide" to the second element's root. - **Ending --ate**: - Used for compounds that contain **oxygen** along with another element. - Indicates the presence of a **polyatomic ion** that includes oxygen. - Typically used in compounds where the nonmetal is combined with oxygen and another element. - Examples: - **NaNO~3~**: Sodium Nitrate (contains the polyatomic ion **nitrate**) - **CaCO~3~**: Calcium Carbonate (contains the polyatomic ion **carbonate**) - The name often reflects the polyatomic ion it contains, with "-ate" indicating a higher oxidation state compared to "-ite," which denotes a lower oxidation state. - **Differences between --ate and --ite**: - **-ate**: Indicates a compound with a higher number of oxygen atoms. - **-ite**: Indicates a compound with fewer oxygen atoms. - Examples: - **NO~3~^-^**: Nitrate (higher oxygen count) - **NO~2~^-^**: Nitrite (lower oxygen count) - **Importance of Understanding Endings**: - Helps in identifying the composition and structure of various compounds. - Allows for better understanding of chemical reactions and properties of substances. - **Summary**: - **-ide**: Binary compounds, usually contain nonmetals. - **-ate**: Compounds with oxygen and polyatomic ions, indicating higher oxidation states. - Recognizing these endings aids in the study of chemistry and the classification of compounds. **Deduction of the formulae of ionic compounds** - **Ionic Compounds**: Formed by the combination of **cations** (positively charged ions) and **anions** (negatively charged ions). - **Key Concepts**: - **Cation**: A positively charged ion, typically a metal. - **Anion**: A negatively charged ion, typically a non-metal or polyatomic ion. - **Charge Balance**: The total positive charge from cations must equal the total negative charge from anions to create a neutral compound. - **Common Types of Ionic Compounds**: - **Oxides**: Contain **O²⁻** ions. - **Hydroxides**: Contain **OH⁻** ions. - **Halides**: Contain halogen ions such as **F⁻**, **Cl⁻**, **Br⁻**, **I⁻**. - **Nitrates**: Contain **NO₃⁻** ions. - **Carbonates**: Contain **CO₃²⁻** ions. - **Sulfates**: Contain **SO₄²⁻** ions. - **Steps to Deduce Formulae**: - **Identify the Ions**: Write down the cation and anion involved. - **Determine Charges**: Note the charges of both ions. - **Cross Multiply**: Use the absolute value of the opposite ion\'s charge as a subscript for the respective ion. - **Simplify**: If needed, reduce the subscripts to the simplest whole number ratio. - **Examples**: - **Sodium Chloride (NaCl)**: - Cation: **Na⁺** (charge +1) - Anion: **Cl⁻** (charge -1) - Formula: NaCl (1:1 ratio) - **Calcium Oxide (CaO)**: - Cation: **Ca²⁺** (charge +2) - Anion: **O²⁻** (charge -2) - Formula: CaO (1:1 ratio)  - **Aluminum Sulfate (Al₂(SO₄) ₃)**: - Cation: **Al³⁺** (charge +3) - Anion: **SO₄²⁻** (charge -2) - Cross multiply: 2 Al³⁺ for every 3 SO₄²⁻ → Al₂(SO₄) ₃ - **Important Note**: Always check for the simplest ratio of ions to ensure the formula is correct. **Structure of an ionic compound as a lattice structure** - **Ionic Compounds**: Formed from the electrostatic attraction between **cations** (positively charged ions) and **anions** (negatively charged ions). - **Lattice Structure**: - An **ionic compound** has a **lattice structure**, which is a three-dimensional arrangement of ions. - Ions are arranged in a **regular, repeating pattern**, creating a strong and stable structure. - **Formation of Lattice**: - When ionic compounds form, cations and anions attract each other due to their opposite charges. - This attraction causes the ions to pack closely together in a specific arrangement. - **Characteristics of the Lattice**: - The **lattice structure** maximizes attraction between oppositely charged ions while minimizing repulsion between like-charged ions. - Each ion is surrounded by ions of opposite charge, creating a strong bond. - **Properties of Ionic Compounds**: - **High Melting and Boiling Points**: The strong forces holding the lattice together require a lot of energy to break. - **Brittleness**: When force is applied, layers can shift, causing like charges to align and repel each other, leading to shattering. - **Electrical Conductivity**: Ionic compounds conduct electricity when dissolved in water or melted, as ions are free to move. - **Examples of Ionic Compounds**: - Common ionic compounds include **sodium chloride (NaCl)** and **magnesium oxide (MgO)**, both of which display the characteristic lattice structure. - **Visual Representation**: Fig.: Lattice Structure of Sodium Chloride  Fig.: Lattice Structure of Magnesium Oxide - **Importance of Lattice Structure**: - The lattice structure is crucial for understanding the physical properties and behavior of ionic compounds in different conditions. - **Summary**: - Ionic compounds are characterized by their **lattice structure**, which is essential for their stability and properties, such as high melting points, brittleness, and electrical conductivity. **Practice Questions** **1. What is an ionic bond?**\ a. A bond formed through electron sharing\ b. A bond formed through electron transfer\ c. A bond formed through nuclear fusion\ d. A bond formed through proton transfer **2. Which of the following best describes a** **cation?**\ a. An atom that has gained electrons\ b. A negatively charged particle\ c. An atom that has lost electrons\ d. A neutral particle **3. In sodium chloride (NaCl), how many electrons does sodium lose?**\ a. Two electrons\ b. Three electrons\ c. No electrons\ d. One electron **4. What property is NOT characteristic of ionic compounds?**\ a. High melting point\ b. Low boiling point\ c. Crystalline structure\ d. Conducts electricity when melted **5. In a calcium ion (Ca²⁺), what happens to the electron count?**\ a. Gains two electrons\ b. Loses two electrons\ c. Gains one electron\ d. No change in electrons **6. What type of elements typically form anions?**\ a. Noble gases\ b. Transition metals\ c. Nonmetals\ d. Alkali metals **7. The ending \"-ate\" in a compound name typically indicates:**\ a. A binary compound\ b. The presence of oxygen\ c. The absence of metal\ d. A reduced state **8. In an ionic lattice structure, what causes the brittleness of ionic compounds?**\ a. Weak bonds between ions\ b. Random arrangement of ions\ c. Like charges aligning when force is applied\ d. High melting points **9. How many electrons does a Group 7 (halogen) element typically gain?**\ a. One electron\ b. Two electrons\ c. Three electrons\ d. Four electrons **10. What is the charge of a sulfate ion (SO₄)?**\ a. -1\ b. +2\ c. -2\ d. +1 **11. The atomic number represents the number of:**\ a. Neutrons in an atom\ b. Protons in an atom\ c. Electrons in an atom\ d. Protons plus neutrons **12. Which group of elements typically forms +2 cations?**\ a. Halogens\ b. Alkaline earth metals\ c. Noble gases\ d. Transition metals **13. In dot and cross diagrams, what do the crosses represent?**\ a. Protons\ b. Neutrons\ c. Valence electrons from one atom\ d. Nuclear particles **14. The \"-ide\" ending in compound names indicates:**\ a. A polyatomic ion\ b. A binary compound\ c. The presence of oxygen\ d. A complex ion **15. What determines the ratio of ions in an ionic compound?**\ a. The atomic mass\ b. The charge balance\ c. The number of neutrons\ d. The electron configuration **Answers** 1. b 2. c 3. d 4. b 5. b 6. c 7. b 8. c 9. a 10. c 11. b 12. b 13. c 14. b 15. b