Chemistry Test Prep PDF

Grade 11 Unit evaluation preparation: Chemical bonding and polarity Shapes ➔ Trigonal pyrimidal There are 4 things surrounding the centeral atom (element) Out of four there is 1 lone pair A lone pair is two dots (: or.. ) on the central atom The angle between each atom is...

Grade 11 Unit evaluation preparation: Chemical bonding and polarity Shapes ➔ Trigonal pyrimidal There are 4 things surrounding the centeral atom (element) Out of four there is 1 lone pair A lone pair is two dots (: or.. ) on the central atom The angle between each atom is 107 ➔ Tetrahydral There are 4 atoms surrounding the central atom (element) The angle between each atom is 109.5 ➔ Trigonal planar There are three atoms surrounding the central atom The angle between each atom is 120 ➔ Bent (1) There are 3 things surround the central atom Out of 3, 1 is a lone pair Angle between each atom is 116  ➔ Bent (2) There are 4 atoms surrounding the central atom Out of 4, 2 are lone pairs The angle between each atom is 105 ➔ Linear There are two atoms surround the central atom There is an angle of 180 between each atom ➔ Molecular Geometry is the shape of a molecule based on the number of bonds (bond pairs) around the central atom ➔ Electron Geometry is the shape of a molecule considering both bond pairs and lone pairs around the central atom. Eg. Let’s look at water (H₂O) as an example for both: Molecular Geometry: Water has two bond pairs (one between each hydrogen and the oxygen). Since only the bond pairs are counted, the molecular geometry is bent or V-shaped. Electron Geometry: Water has two bond pairs and two lone pairs (on the oxygen atom). When counting both bond and lone pairs, the electron geometry is tetrahedral, because there are four regions of electron density (2 bond pairs + 2 lone pairs). Key : The bond pairs tend to push the other atoms down due to their repulsion. Steps to figuring it out: 1. Draw the lewis dot diagram for it 2. Consider moloecular geometry first ( only the bond pairs)  3. Consider the electron geometry ( bond & lone pairs) Bond pairs Lone pairs Molecular geometry Electron geometry 1 0 linear linear 2 0 Linear linear 3 0 Trigonal planar Trigonal planar 2 1 Bent Trigonal planar 4 0 Tetrahedral Tetrahedral 3 1 Trigonal pyramidal Tetrahedral 2 2 Bent Tetrahedral 1 3 Linear Tetrahedral Bond pair: Electrons shared between two atoms, forming a bond (e.g., in H₂O, the hydrogen and oxygen atoms are connected by bond pairs). (Simple words: Bond pair is the line that connects two atoms when it is bonding ) Lone pair: Electrons that stay on a single atom without being shared (e.g., in H₂O, oxygen has two lone pairs of electrons that do not bond with hydrogen). (Simple words: A lone pair is the two dots which act as a line but are not connected, when we look at determining the shape we only consider the lone pairs which are on the central atom) Lab - Background 1. Ionic compounds are formed when oppositely charged ions (positive and negative) attract each other. This attraction is called an ionic bond. These compounds usually form crystals where the ions are arranged in a regular pattern. 2. Covalent compounds are held together when atoms share electrons. These form molecules, which are groups of atoms bonded together. The molecules themselves stick together in clusters, but the forces holding them together (called intermolecular forces) are weaker than the forces in ionic compounds. These bonds affect properties like: Melting Point: Ionic compounds need a lot of heat to melt because their bonds are strong. Covalent compounds have weaker forces between molecules, so they melt at lower temperatures.  Volatility: Volatile substances have weaker bonds, allowing particles to escape easily and be smelled. Solubility: Ionic compounds tend to dissolve easily in water because water can pull apart the ions. Covalent compounds usually don’t dissolve as well in water. ○ If we take sugar and salt as an example, sugar is a covalent bond meaning its non-metal + non metal. When you put this into water it will break into individual molecules that make it up making it harder to dissolve. Salt is an ionic bond meaning non metal + metal. When you put it into water it breaks up into individual atoms causing it to dissolve way faster. Conductivity: When a compound dissolves in water, it might conduct electricity if the dissolved particles can move freely. If it conducts, the compound is called an electrolyte and likely has ionic bonds. If it doesn’t, it’s called a non-electrolyte and likely has covalent bonds Trends ➔ Ionization energy: Energy required to remove an electron from an atom - Decreases down a period due to weaker electronegativity and larger atomic radius - Increases across a period due to stronger electronegativity and smaller atomic radius ➔ Atomic radius: the size of an atom - Increases down a period due to weaker electronegativity - Decreases across a period due to stronger electronegativity ➔ Proton pull (nuclear attraction): the ability of the protons in the nucleus to pull the electrons closer towards the nucleus. - Increases across a period due to weaker shielding effect - Decreases down a period due to stronger shielding effect ➔ Shielding effect: when the innermost electrons “block” or “shield” the outermost electrons from the proton pull - Decreases across a period due to stronger proton pull - Increases down a period due to weaker proton pull (add test questions) Polarity Polar Bond: A polar bond occurs when two atoms share electrons unequally. One atom has a stronger pull on the electrons (higher electronegativity), causing one end of the bond to be slightly negative and the other end to be slightly positive. This creates a dipole (a separation of charge). Example: In water (H₂O), the oxygen atom pulls the shared electrons more than the hydrogen atoms, making the bond polar. Nonpolar Bond:  A nonpolar bond happens when two atoms share electrons equally. This usually occurs between identical atoms or when the atoms have similar electronegativities, so there’s no charge separation. Example: In oxygen gas (O₂), the two oxygen atoms share electrons equally, making the bond nonpolar. To determine whether a bond is polar or nonpolar, you can follow these steps: 1. Find Electronegativity Values: Look up the electronegativity values of the two atoms involved in the bond. Electronegativity is a measure of an atom's ability to attract and hold onto electrons. You can use a table of electronegativity values 2. Calculate the Difference: Subtract the electronegativity of one atom from the other to find the difference 3. Determine Polarity Based on the Difference: Use the following guidelines for the electronegativity difference: ○ Nonpolar Covalent Bond: If the difference is 0 (e.g., between identical atoms, like Cl-Cl or H-H). ○ Polar Covalent Bond: If the difference is between 0.1 and 1.7 (e.g., H-Cl). ○ Ionic Bond: If the difference is greater than 1.7 (e.g., Na-Cl). Example: Let’s say we want to determine the polarity of the bond between hydrogen (H) and chlorine (Cl): 1. Electronegativity Values: ○ H = 2.1 ○ Cl = 3.0 2. Calculate the Difference: Difference= 3.0−2.1 = 0.9 3. Determine Polarity: ○ Since the difference (0.9) is between 0.1 and 1.7, the H-Cl bond is polar. Polar Molecule Polar molecules Non-polar molecules When the negative charge is unevenly When the negative charge is evenly distributed around the atom distributed around the atom Asymmetrical shape (bent or pyramidal) Symmetrical shape (linear, tetrahedral or  At least 1 polar bond trigonal planar) Polar bonds Intermolecular Forces Intermolecular forces are the forces of attraction or repulsion between molecules. These forces are responsible for the physical properties of substances, such as boiling points, melting points, and solubility. Here are the main types of intermolecular forces: 1. Dipole-Dipole Forces: Occur between polar molecules, where the positive end of one molecule is attracted to the negative end of another. Stronger than London dispersion forces but weaker than hydrogen bonds. 2. Hydrogen Bonds: A special type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like nitrogen (N), oxygen (O), or fluorine (F). Hydrogen bonds are stronger than typical dipole-dipole interactions and play a crucial role in the properties of water and biological molecules like DNA. 3. London Dispersion Forces Present in all molecules, whether polar or nonpolar, but are the only forces acting in nonpolar molecules. Caused by temporary shifts in electron density that create instantaneous dipoles. Generally weaker than dipole-dipole forces and hydrogen bonds. 4. Ion-Dipole Forces: Occur between an ion and a polar molecule. This force is particularly important in solutions, such as when ionic compounds dissolve in water. Intermolecular forces help explain why some substances are gases, while others are liquids or solids at room temperature, and influence properties like boiling point, melting point, and solubility. 1. Kelsi vs. Iesha on Ammonia (NH₃) (CHATGPT RESPONSE TO HELP FOR TEST) Kelsi is correct; ammonia (NH₃) is a polar molecule. Reasoning: ○ Shape: Ammonia has a trigonal pyramidal shape due to the presence of one lone pair on the nitrogen atom. While Iesha refers to a trigonal planar shape, that applies to a molecule like boron trifluoride (BF₃), which has no lone pairs and is nonpolar. In NH₃, the lone pair creates an uneven distribution of electron density. ○ Electronegativity: The nitrogen atom is more electronegative than hydrogen atoms, resulting in a partial negative charge near the nitrogen and partial positive charges near the hydrogens. ○ Overall Dipole Moment: The shape and the electronegativity differences create a net dipole moment, making NH₃ polar. 2. Predicting the Highest Boiling Point a) CH₃OH and CH₃Cl Answer: CH₃OH (methanol) will have the higher boiling point. Reason: Methanol can form hydrogen bonds due to the presence of the -OH (hydroxyl) group, which are strong intermolecular forces. CH₃Cl (methyl chloride) only has dipole-dipole interactions and London dispersion forces, which are weaker than hydrogen bonds. b) H₂O and H₂S Answer: H₂O (water) will have the higher boiling point. Reason: Water has strong hydrogen bonds because of the highly electronegative oxygen atom and the presence of hydrogen atoms. H₂S (hydrogen sulfide) has weaker dipole-dipole interactions and London dispersion forces, resulting in a lower boiling point. c) CH₄ and C₃H₈ Answer: C₃H₈ (propane) will have the higher boiling point. Reason: Propane has a larger molecular size and more surface area than methane (CH₄), leading to greater London dispersion forces. Although both are nonpolar, the larger number of electrons in C₃H₈ allows for stronger temporary dipoles, which contribute to a higher boiling point compared to CH₄. Drawing Lewis structure Elements tha tdon’t require a full octet Hydrogen and Helium: Stable with 2 electrons. Beryllium: Stable with 4 electrons. Boron: Stable with 6 electrons. Third Period Elements (P, S, Cl): Can have expanded octets, accommodating more than 8 electrons. Steps 1. Count Valence Electrons: a. Find the number of valence electrons for each atom. Add them up for the entire molecule. 2. Choose a Central Atom: a. Usually, it's the least electronegative atom (not hydrogen). 3. Draw a Skeleton Structure: a. Connect atoms with single bonds. (Example: A central atom connected to surrounding atoms.) 4. Distribute Remaining Electrons: a. Place leftover electrons on the outer atoms first to complete their octets (or duets for hydrogen). 5. Complete the Central Atom's Octet: a. If the central atom doesn’t have an octet, form double or triple bonds by sharing lone pairs from outer atoms. 6. Check Your Structure: a. Make sure all atoms have the correct number of electrons and the total matches the initial count. Example Steps to Draw the Lewis Structure for Phosphoric Acid (H₃PO₄) 1. Count Valence Electrons: ○ Hydrogen (H): 1 valence electron × 3 = 3 ○ Phosphorus (P): 5 valence electrons (Group 15) ○ Oxygen (O): 6 valence electrons × 4 = 24 ○ Total: 3 (H) + 5 (P) + 24 (O) = 32 valence electrons 2. Choose a Central Atom: ○ Phosphorus (P) is the central atom because it can form multiple bonds and is less electronegative than oxygen. 3. Draw a Skeleton Structure: ○ Connect phosphorus to four oxygen atoms. 4. Distribute Remaining Electrons:  ○ After using 6 electrons for the 3 O-H bonds (2 for each), we have 26 electrons left (32 - 6 = 26). ○ Distribute electrons among the oxygen atoms to complete their octets. Start by placing 6 electrons (3 lone pairs) on each of the three oxygens connected to hydrogen. 5. Complete the Central Atom's Octet: ○ At this point, the phosphorus atom still needs to be connected properly. ○ The last oxygen is connected to phosphorus with a single bond and needs more electrons. Move one of the lone pairs from the adjacent oxygen atom (the one bonded to hydrogen) to form a double bond between phosphorus and that oxygen atom. Only elements that can form multiple covalent bonds with itself or eachother are… Carbon Nitrogen Oxygen Sulfur Phosphorus