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CompatibleCognition2761

Uploaded by CompatibleCognition2761

Manal Alzahmi

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chemistry periodic table atomic structure ionic bonding

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This document contains questions and explanations in the chemistry field, likely for revision or study. The content includes topics like predicting periodic properties of elements, understanding quantum numbers, electron configurations, and forming ionic bonds.

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Chem revision – Based on MOE coverage: Q1: Predict the periodic properties of elements (e.g: atomic radius, ionization energy, and electronegativity) in the period and group in the periodic table. Atomic Radius:  Across a Period (→): Decreases ➔ More protons pull electrons closer to th...

Chem revision – Based on MOE coverage: Q1: Predict the periodic properties of elements (e.g: atomic radius, ionization energy, and electronegativity) in the period and group in the periodic table. Atomic Radius:  Across a Period (→): Decreases ➔ More protons pull electrons closer to the nucleus.  Down a Group (↓): Increases ➔ More electron shells are added, increasing size. Ionic Radius:  Cations (positive ions): Smaller than their parent atoms. ➔ Losing electrons reduces electron-electron repulsion.  Anions (negative ions): Larger than their parent atoms. ➔ Gaining electrons increases repulsion, expanding the ion.  Across a Period (→): Decreases (for ions of the same charge). ➔ Higher nuclear charge pulls electrons closer.  Down a Group (↓): Increases ➔ More electron shells increase size. Ionization Energy (IE):  Across a Period (→): Increases ➔ Electrons are held more tightly due to higher nuclear charge.  Down a Group (↓): Decreases ➔ Electrons are farther from the nucleus and easier to remove. Electronegativity:  Across a Period (→): Increases ➔ Atoms want to complete their octet (non-metals attract electrons strongly).  Down a Group (↓): Decreases ➔ Larger atoms have weaker attraction for bonding electrons. Q2: Identify the four quantum numbers and their respective significance and calculate the number of electrons in each level. Principal Quantum Number (n), What does it tell us:  The energy level or shell of an electron.  How far the electron is from the nucleus. Energy level (n): Subshells: Number of orbitals: Max electrons: n=1 s 1 2 n=2 s,p 1+3=4 8 n=3 s,p,d 1+3+5=9 18 n=4 s,p,d,f 1+3+5+7=16 32 Each energy level (n) can hold 2𝑛2 electrons. Q3 + Q4: Write the electronic configuration of a variety of elements of the periodic table, employing the Pauli exclusion principle, the Hund rule, and the Aufbau principle for upward building. Full - Electron configuration: 1𝑠 2 2𝑠 2 2𝑝6 3𝑠 2 3𝑝6 4𝑠 2 3𝑑10 4𝑝 6 5𝑠 2 4𝑑10 5𝑝6 6𝑠 2 4𝑓14 5𝑑10 6𝑝6 7𝑠 2. Aufbau Principle:  Electrons fill the lowest energy orbitals first. Pauli Exclusion Principle:  Only 2 electrons per orbital, with opposite spins. Hund’s Rule:  Electrons spread out in orbitals of the same energy before pairing up. Q5: State the key features of the periodic table. Periods: Horizontal rows (7 in total). Groups: Vertical columns (18 in total). Metals, Non-metals, Metalloids: Classified based on properties. Key Groups: Alkali metals (Group 1), Alkaline earth metals (Group 2), Halogens (Group 17) and Noble gases (Group 18). Representative elements (Group 1,2 and 13-18) and transition elements (Group 3-12). Inner transition elements (Lanthanide series, actinide series). Trends: Atomic size, ionization energy, and electronegativity show predictable patterns. Q6: Describe the different groups of elements in the periodic table and predict the physical and chemical properties of each group. Go back to the book for this question, I brought the table from the well know ChatGPT. Q7: Predict the periodic properties of elements (e.g: atomic radius, ionization energy, and electronegativity) in the period and group in the periodic table.  Based on your knowledge of periodic trends. Q8: Illustrate how do positive and negative ions form. Positive ions (cations) are formed when atoms (metals) lose electrons. Negative ions (anions) are formed when atoms (non-metals) gain electrons Q9: Illustrate how ionic bonds are formed and how ions are arranged in an ionic compound. 1. Electron transfer occurs between a metal (loses electrons) and a non-metal (gains electrons). 2. The metal becomes a positive ion (cation), and the non-metal becomes a negative ion (anion). 3. The opposite charges attract, forming an ionic bond. 4. The ions arrange themselves in a crystal lattice to form a stable ionic compound. A crystal lattice in ionic compounds is a repeating 3D structure of ions arranged in a regular pattern, where opposite charges attract. This structure gives ionic compounds their strength and stability. Example of Ionic Bond: - Na (sodium) + Cl (chlorine) → NaCl (sodium chloride) - This bond forms because of the electrostatic attraction between Na⁺ and Cl⁻. Q10: Explain the structure and properties of ionic compounds based on their bond types, strength and organization. Ionic bonds form between metals (positive ions) and non-metals (negative ions), creating a strong bond due to electrostatic attraction. The ions arrange themselves in a crystal lattice, forming a stable, repeating 3D structure. Ionic compounds have strong bonds, making them hard, brittle, and having high melting and boiling points. As called “electrolyte” they conduct electricity when melted or dissolved in water, as ions can move freely in these conditions. Q11: Write chemical formulas for binary and polyatomic compounds or use simulation software to show them, including those that have more than one oxidation number, naming it by using nomenclature system of International Union of Pure and Applied Chemistry (IUPAC). 1. Identify the ions involved (including their oxidation states). 2. Write the symbols of the ions and balance the charges to make the compound neutral. 3. For metals with more than one oxidation state, use Roman numerals in the name. 4. For polyatomic ions, use the common names of the ions. A. Binary Compounds with Metals and Non-metals: 1. Write the formula by combining the ions of the two elements. 2. Balance the charges to ensure the compound is neutral (no overall charge). 3. Name the metal first, followed by the non-metal with an ending of -ide. B. Binary Compounds with More Than One Oxidation State (Metal + Non-metal): Example: Iron(II) Chloride (FeCl₂) and Iron(III) Chloride (FeCl₂)  Iron (Fe) can form two ions: Fe²⁺ (oxidation state +2) or Fe³⁺ (oxidation state +3).  Chlorine (Cl) forms a Cl⁻ ion. Iron(II) Chloride (FeCl₂):  Fe²⁺ needs two Cl⁻ ions to balance the charges, so the formula is FeCl₂. Iron(III) Chloride (FeCl₂):  Fe³⁺ needs three Cl⁻ ions to balance the charges, so the formula is FeCl₂. C. Polyatomic Compounds: Polyatomic compounds involve ions that are made of more than one atom. These are usually made of non-metals and contain polyatomic ions like sulfate (SO₂²⁻), nitrate (NO₂⁻), ammonium (NH₂⁺), etc. Example: Calcium Carbonate (CaCO₃ )  Calcium (Ca) forms a Ca²⁺ ion.  Carbonate (CO₃ ²⁻) has a charge of -2.  The charges balance out (1 calcium ion and 1 carbonate ion), so the formula is CaCO₃. Q12: Explain the structure and properties of metallic compounds based on their bonds types, strength and organization.  Metallic bonds occur when metal atoms share their valence electrons in a sea of electrons.  The atoms are arranged in an organized lattice that allows them to slide past each other.  Metals are conductive, malleable, ductile, strong, and shiny because of the unique bonding and structure of metal atoms. In simple terms, metallic compounds are strong because of the sea of electrons that hold the atoms together, and they have great properties like conductivity and flexibility because those electrons can move freely. Arranged in a Lattice: Metal atoms are arranged in a regular, repeating pattern called a lattice. This arrangement is very orderly, but the atoms are still able to slide past each other easily because of the sea of electrons. Why the Lattice is Important: The organized structure allows metals to be ductile (can be stretched into wires) and malleable (can be hammered into sheets), as the atoms can slide past each other without breaking the bonds. Bond Strength: The strength of the metallic bond depends on two factors:  Number of valence electrons: More electrons in the sea of electrons usually means a stronger bond. For example, metals like copper and gold have fewer valence electrons, while tungsten and iron have more, leading to stronger metallic bonds.  Size of the metal atoms: Smaller atoms can pack together more tightly, leading to stronger metallic bonds. Effect on Properties: Metals with stronger metallic bonds tend to be stronger and have higher melting points. Properties:  High electrical and thermal conductivity.  Malleability (Hammered into sheets) and ductile (Stretched to wires).  High melting point but extreme boiling point (Boling point is higher than the melting point).  Metals have great strength because of the strong attraction between metal atoms and the sea of electrons. Q13: Define both sigma and pi bonds and their differences. A sigma bond is the strongest type of covalent bond. It forms when two atomic orbitals (like an s orbital or a p orbital) overlap head-on.  This bond form between two s orbitals (ss), s and a p orbital (sp), or two p orbitals (pp). A pi bond forms when two p orbitals overlap side-by-side, not head-on.  Pi bonds only form after a sigma bond has already been made (e.g., in double or triple bonds).  Weaker than sigma bonds because the sideways overlap is not as strong as the head- on overlap of a sigma bond. Sigma bonds are the primary bond that holds atoms together, while pi bonds are additional bonds that provide extra strength in double or triple bonds but are weaker than sigma bonds. Q14: Explains the relationship between strength, length, and dissociation energy of a covalent bond.  Stronger bonds → Shorter bond length → Higher dissociation energy.  Weaker bonds → Longer bond length → Lower dissociation energy. Bonds and dissociation energy have a direct relationship, while length and both bonds and dissociation energy have an inverse relationship (opposite). Q15: Name a binary molecular compound from its molecular formula. 1. Identify the two elements: Look at the two different elements in the formula. The first element is usually non-metal, and the second element is also non-metal. 2. Use prefixes for the number of atoms: Use prefixes to indicate how many atoms of each element are present. 3. Name the first element: The first element is named as is (no change to its name). 4. Name the second element: The second element is named with an -ide ending (like "chloride," "oxide"). Note that only and only if we had a one atom for the first element, we should NOT write the prefix “mono”. We are starting from 2 and above. Q16: Name acidic solutions. 1. Naming Binary Acids (Acids with Two Elements): When the acid consists of hydrogen (H) and a non-metal (such as chlorine, sulfur, etc.), follow these steps: 1. Start with the prefix "hydro-" to indicate it is acidic. 2. Name the non-metal (second element) and change its ending to -ic. 3. Add the word "acid" at the end. 2. Naming Oxyacids (Acids with Hydrogen and Oxygen): When the acid consists of hydrogen (H), oxygen (O), and a non-metal or polyatomic ion (such as sulfate or nitrate), follow these steps: 1. Identify the polyatomic ion and look at its ending. o If it ends in -ate, the acid name will end in -ic. o If it ends in -ite, the acid name will end in -ous. 2. Add the word "acid" at the end. Q17: Represents molecules that are exceptions to the octet rule and explan these exceptions Incomplete Octet (Less than 8 electrons):  Elements like boron (BF₂) and beryllium (BeCl₂) can form stable compounds with fewer than 8 electrons in their valence shell. Expanded Octet (More than 8 electrons):  Elements in Period 3 and beyond (such as phosphorus (PCl₂) and sulfur (SF₂)) can have more than 8 electrons in their valence shell due to the availability of d orbitals. Odd Number of Electrons (Free Radicals):  Some molecules like NO₂ have an odd number of electrons, meaning it's impossible to satisfy the octet rule for all atoms in the molecule. Q18: Represents molecules that are exceptions to the octet rule and explain these exceptions Steps to Draw a Lewis Structure:  Step 1: Count Valence Electrons: o Look at the periodic table to find the valence electrons for each atom.  Step 2: Identify the Central Atom: o The least electronegative atom (except hydrogen) is usually the central atom. o Hydrogen is ALWAYS at the side, never at the center.  Step 3: Draw Single Bonds: o Connect the atoms using single bonds (each bond uses 2 electrons).  Step 4: Distribute Remaining Electrons: o Place the remaining electrons as lone pairs around the outer atoms. o Make sure each atom (except hydrogen) gets an octet (8 electrons).  Step 5: Make Double or Triple Bonds (if needed): o If an atom doesn’t have 8 electrons, move lone pairs to form double or triple bonds. Q19: Represents molecules that are exceptions to the octet rule and explain these exceptions.  You should memorize it from these pages (117 and 118). Q20: Identify the differences and similarities between polar and nonpolar covalent bonds and polar and nonpolar molecules. Covalent bonds form when two atoms share electrons. The difference between polar and nonpolar covalent bonds depends on how equally the electrons are shared between the atoms. Polar Covalent Bond:  Electrons are shared unevenly.  This happens when one atom is more electronegative (stronger attraction for electrons) than the other, causing the electrons to spend more time closer to the more electronegative atom.  This creates a partial negative charge (δ-) on the more electronegative atom and a partial positive charge (δ+) on the other atom. Polar Molecule:  The molecule has a positive and a negative side (a dipole).  This happens if the molecule has polar covalent bonds, and the shape doesn’t cancel out the dipoles.  The molecule is asymmetrical, so the dipoles do not cancel each other out. Nonpolar Covalent Bond:  Electrons are shared evenly.  This happens when two atoms have the same electronegativity or their difference is very small, so they share the electrons equally.  There are no partial charges (δ+ or δ-) because the sharing is equal. Nonpolar Molecule:  The molecule has no overall charge or dipole.  This happens if the molecule either has nonpolar covalent bonds or polar covalent bonds that cancel out due to symmetry (e.g., the molecule is symmetrical). Electronegativity difference: Bond character: >1.7 Mostly ionic bond. 0.4 – 1.7 Polar covalent bond.

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