Chem Review Unit Test #1.pdf

Janujan Vaseekaran Lesson 1: Unit 1  Janujan Vaseekaran Lesson 2: Unit 1 - The Quantum mechanical Model Sublevels - S , P , D , F - Sublevels further divided into orbitals - Orbitals: a...

Janujan Vaseekaran Lesson 1: Unit 1  Janujan Vaseekaran Lesson 2: Unit 1 - The Quantum mechanical Model Sublevels - S , P , D , F - Sublevels further divided into orbitals - Orbitals: a region of space where there is a high probability of finding an electron  Janujan Vaseekaran - S sublevel a maximum of 2 electrons - P sublevel a maximum of 6 electrons - D sublevel a maximum of 10 electrons - F sublevel a maximum of 14 electrons Orbital diagrams - included the energy levels, sublevels and orbitals - Electrons must be put from lowest to highest energy as with bohr rutherford diagrams - the 1st energy level has less energy than the 2nd energy level which has energy that the 3rd Rules to guide the placing of electrons in an orbital diagram: - Aufbau principle: - Aufbau in german means building up - Each electron added to the lowest energy level available - Pauli Exclusion Principle - A maximum of 2 electrons can occupy any one orbital - Hund's rule: - One electron occupies each of several orbitals at the same energy before a second electron can occupy the same orbital Electron Configurations: Unit 1 Lesson 2b - The Quantum Mechanical Model (continued)  Janujan Vaseekaran Anions - negative ions - Add a extra electron and then distribute them Cation - positive ions - Must draw orbital diagram for neutral atom firs - Electron must be removed from the highest energy level Condensed Electron Configuration: - Some of the electrons of an atom are represented using a symbol of the preceding noble has - Only the remaining electrons beyond the nobel gas are shown - Because only the outer electrons are important for explaining the chemical properties Multiple Ion Charges: - Larger atom ---> more complex electron configuration - More overlapping of outer energy levels - Different options for electrons to be removed  Janujan Vaseekaran Lesson 3 - Unit 1 - Quantum Numbers - Bohrs model explained many things about small atoms but later failed to explain the behaviour of atoms with more than one electron it was later found that electrons do not really travel in fixed paths - Heisenberg Uncertainty Principle - states that position and velocity of an electron cannot be measured at the same time because an electron is so small - De Broglie - Wave particle duality - Electrons have many characteristics of waves, electrons behave like particles and other times like waves - Schrodinger - The motion and energy of electrons are described by extremely complexed equations - To find maximum of orbitals in energy level = n^2 - To find maximum of electrons in energy level 2n^2 Sigma & Pi Bonds: Covalent Bonds: - Form between two nonmetals  Janujan Vaseekaran - Form between p and/or s orbitals - May be single, double or triple bonds Examples of Single Bonds: Lone Pair: two electrons in outer energy level that do not take part in bonding Sigma bond: - Cylindrical symmetrical about the axis joining the 2 nuclei - Formed by a head-on overlap of p or s orbitals - A single bond or the first bond in a double or triple bond Pi Bond: - Electron density divided into 2 regions (create high electron density) - Above and below the axis joining the two nuclei - Formed by a parallel overlap of two p orbitals - 2nd in a double bond - 2nd or 3rd bond in a triple bond - Pi bond Less stable than a sigma bond - Because of the extent of overlapping with sigma bonds it has more overlapping which makes it more stable compared to less overlapping of the pi bond  Janujan Vaseekaran Molecular Shape: Covalent Bonding: - Bonds form to make atoms more stable - In general atoms like to complete an outer energy level where they bond - Covalent bonds are formed when electrons are shared between two atoms - Covalent bond exist in solids when electrons are shared between two atoms - Covalent bonds exist in solids, liquids and gases Lewis structures:  Janujan Vaseekaran Incomplete and Expanded Octets: Incomplete Octet: Expanded Octet: Orbital diagrams work well for describing individual, isolated atoms. However, when atoms bond, hybridization occurs, allowing atoms like carbon to form more bonds than you would predict from the basic orbital diagram. In the case of carbon, it hybridizes to form four bonds, even though the ground-state orbital diagram suggests it has only two unpaired electrons. VSEPER theory - Valence Shell Electron Pair Repulsion Theory: - Only valence electrons of the central atom are involved in determining a molecule’s shape - Electrons repel each other and therefore position themselves as far apart as possible - Lone pairs electrons are less bound to the nucleus and therefore take up more space  Janujan Vaseekaran Hybridization Theory: The text describes Hybridization Theory, which explains how atomic orbitals mix or "hybridize" to form new orbitals during the formation of covalent bonds around a central atom. These new orbitals are of intermediate energy compared to the original orbitals. The theory aligns with VSEPR (Valence Shell Electron Pair Repulsion) theory, which states that these hybridized orbitals will arrange themselves as far apart as possible in three-dimensional space to minimize repulsion between electron pairs. This arrangement helps to explain the molecular geometry and bonding behaviors that cannot be fully predicted by simple orbital diagrams of isolated atoms.  Janujan Vaseekaran Polarity: Electronegativity: the tendency to attract electrons within a bond - The Pulling power that an atom has on shared electrons Bond Polarity: a slight difference in charge between bonded atoms which have different electronegativity values Molecular polarity: - A difference in charge between sides of an entire molecule - Just because a molecule contains polar bonds does not necessarily mean it will be a polar molecule Intermolecular Forces: Intramolecular Bonds: - The force that holds atoms together within a molecule - Breaking & Forming intramolecular bonds are chemical change  Janujan Vaseekaran Intermolecular Forces: - The force that exists between the molecules - Interaction between particles - Gas particles are so far apart that there is very little interaction between molecules - In solids, & liquids the particles are close enough to interact. In these states intermolecular forces become significant. - Breaking & forming intermolecular forces are physical changes Three common types: - DIspersion Forces - Dipole Dipole Forces - Hydrogen Bonding # the stronger the intermolecular forces, the higher the melting point and boiling point of the substance 1. Dispersion - Weak attractive forces between Nonpolar atoms or molecules - Electrons in both molecules are randomly distributed - Electrons in one molecule shift to one side momentarily - A temporary dipole or change forms - When the non-polar covalent molecule has its electrons close and makes the electron distribution uneven which can cause one side of the atom to become more negatively charged thus creating temporary dipole - The temporary positive charge on one end of a molecule induces electrons to shift in a nearby molecule and attraction between the molecules results Dispersion forces are stronger when there are more electrons in the molecule Dispersion forces are stronger when the molecule has a larger surface area Dipole Dipole forces - Strong forces of attraction between two or more polar molecules - The positive and negatively charged ends of each molecule attract - The charges on polar molecules are permanent and strong. Polar molecules tend to have higher boiling point than nonpolar molecules.  Janujan Vaseekaran #factors that affect dipole dipole forces are the number of electrons in the molecule and the difference in electronegativity between the atoms Hydrogen bonding: - every strong form of dipole dipole attraction - Occurs when a molecule has a hydrogen atom bonded to an oxygen, nitrogen, and fluorine atom - The positive dipole of the oxygen atom on a nearby molecule this is called a hydrogen bond Hydrogen bonding is responsible for several of the properties of water: - Relatively high boiling point of water - High surface tension - Solid water is less dense than liquid water  Janujan Vaseekaran Lesson 6 - Unit 1 - Models of Chemical Bonding Covalent (Molecular) Bonding (pg 219) Types of Covalent Bonds: 1. Polar Covalent Bonds: Electrons are not shared equally between atoms; they spend more time near the more electronegative atom. 2. Non-Polar Covalent Bonds: Electrons are shared almost equally between the bonded atoms. Intermolecular Forces The three intermolecular forces that hold molecular solids together are: 1. Dispersion Forces (also known as London dispersion forces) 2. Dipole-Dipole Interactions 3. Hydrogen Bonding Metallic Bonding (pg 211-214) Characteristics of Metal Atoms: Metal atoms generally have low electronegativities and their valence shells are less than half-filled. This allows the valence electrons to be delocalized over the metal lattice, which gives rise to metallic bonding. Melting and Boiling Points: Determined by the strength of the metallic bonds, which depend on the number of delocalized electrons and the size of the metal ions. Trends in Metal Melting Points: 1. Trend 1: Melting points generally increase across a period due to increasing number of delocalized electrons. 2. Trend 2: Melting points decrease down a group as atomic size increases, reducing the strength of the metallic bond.  Janujan Vaseekaran Electrical and Thermal Conductivity: Metals are good conductors because their delocalized electrons can move freely through the lattice, allowing them to carry charge (electricity) and transfer thermal energy. Malleability and Ductility: Metals are malleable and ductile because the layers of atoms in a metal can slide over one another without breaking the metallic bonds. Hardness: Hardness is a measure of how resistant a material is to deformation. In metals, hardness increases with stronger metallic bonds. Harder metals: Metals with more delocalized electrons and smaller atoms tend to be harder. Softer metals: Can be made softer by annealing, which allows the metal atoms to rearrange more easily. Alloys: An alloy is a mixture of metals or a metal with another element. Alloys are made to enhance certain properties. Example of Alloy: Steel (an alloy of iron and carbon) is used in construction due to its strength. Substitutional vs. Interstitial Alloys: Substitutional Alloys: Atoms of one metal replace atoms of another metal in the lattice. Interstitial Alloys: Smaller atoms fill in the gaps (interstices) between the metal atoms in the lattice. Ionic Bonding (pg 214-218) Characteristics of Elements in Ionic Bonds: Typically involve a metal (low electronegativity) and a non-metal (high electronegativity). The metal donates electrons to the non-metal, creating oppositely charged ions that form a lattice. Melting and Boiling Points: Determined by the strength of the ionic bonds; stronger ionic bonds result in higher melting and boiling points. For example, sodium chloride (NaCl) has a high melting point because of its strong ionic bonds. Solubility in Water: Ionic compounds are soluble in water when the ionic bonds can be broken by the polar water molecules, allowing the ions to dissolve. Brittleness: Ionic compounds are brittle because when the lattice is disrupted (e.g., by a force), like charges may align and repel, causing the material to break. Electrical Conductivity:  Janujan Vaseekaran Ionic compounds can conduct electricity when molten or dissolved in water, as the ions are free to move and carry charge. Network Solids (pg 224-225) Network Solids: Solids where atoms are connected by a continuous network of covalent bonds throughout the material. Allotropes of Carbon: Allotropes are different structural forms of the same element. Four allotropes of carbon include: 1. Diamond 2. Graphite 3. Fullerene 4. Carbon nanotubes Graphite: Slipperiness: Graphite is slippery because its layers can slide over each other easily due to weak van der Waals forces between them. Diamond: Hardness: Diamond is extremely hard because each carbon atom forms four strong covalent bonds to other carbon atoms, creating a very rigid structure. Lesson 6 - Models of Chemical Bonding Janujan Vaseekaran Janujan Vaseekaran