SCH3U Unit 1 Past Paper PDF

Emission Spectra Demo Back to Bohr. Watch as I demonstrate how distinct colours are visible when elements are excited. These colours are consistent for that element, since electrons are getting excited and then relaxing to exact positions on energy levels - it’s the SAME amount of energy each time...

Emission Spectra Demo Back to Bohr. Watch as I demonstrate how distinct colours are visible when elements are excited. These colours are consistent for that element, since electrons are getting excited and then relaxing to exact positions on energy levels - it’s the SAME amount of energy each time (therefore same photon frequency) BUT… electrons don’t fill perfectly level by level. The quantum model explains why, and you can understand this by looking at subshell electron configurations (extension, relevant in L4/L5). The filling order you must have memorized for this year is: 2 8 8 2 …since you’ll only be asked to draw B-R diagrams up to calcium (atomic # 20) Isotopes We will dig into these in a separate class again with an activity, but it’s worthwhile to understand them for the sake of understanding mass number. Isotopes are different versions of the same element. What is unique about them is the number of neutrons. Isotopes So, for example, a sample of lithium would consist of both of these atoms - some of which have 3 neutrons (lithium-6) and some of which have 4 (lithium-7) You should know which isotope you’re working with when drawing a B-R diagram! Bohr-Rutherford Diagrams Step 1: Identify the atomic number and mass number of the element Step 2: Subtract the atomic number from the mass number to determine the number of neutrons -16 Step 3: State the number of protons and neutrons in the nucleus using the notation seen to the right → Step 4: Begin filling electrons level by level with the limits 2, 8, 8, 2 for each level Key Reminders Step 5: Electrons must be paired only after each location Notation for nucleus in a level has one electron Pairing behaviour Bohr-Rutherford Diagrams Step 1: Let’s draw lithium-7 together Step 2: Let’s draw carbon-12 together Step 3: Let’s draw the entire periodic table up to calcium on the board Step 4 (HW): I will assign a diagram for you to draw and submit independently Key Reminders Notation for nucleus Pairing behaviour  Outermost energy level. Exposed and The Octet Rule important for bonding Atoms tend to want a stable “octet” (8) of electrons in their valence level. This means that elements that are close to having a full valence level will want to have additional electrons to fill it up, and elements that are close to having an empty valence level will want to lose them so that the previous level, which is complete, becomes the outermost one. This change in # of electrons is accomplished in two ways: ions and covalent bonding Anions In order to have a stable valence level (octet), would oxygen be more likely to lose 6 electrons (empty completely), or gain 2 electrons (fill up)?  Bonus: what is the issue with the placement of electrons in this diagram based on my Anions recommended steps? 8p+ - 8p+ + 2e → 8n0 8n0 It is way less energetically demanding to gain 2 electrons. Take a look at the notation for the ion. Cations + In order to have a stable valence level (octet), 0 would sodium be more likely to lose 1 electron (empty completely), or gain 7 electrons (fill up)? Cations 11p+ - - 1e → 11p+ 12n0 12n0 Careful! LOSING an electron makes the ion positive (like subtracting a It takes WAY less energy to lose 1 electron negative number)  Isoelectronic Series 8p+ 11p+ 8n0 neon atom 12n0 (noble gas) oxide ion sodium ion An isoelectronic series all has the same electron arrangement. This emphasizes that neutral atoms tend to become ions in order to stabilize their valence level, as is the case with the very unreactive noble gases. Isoelectronic Series We will return to this topic when we discuss: Trends in atomic/ionic radius within this unit… Stay tuned! This is standardized as having a mass of 12u - internationally agreed upon. “Relative Atomic Mass” is just the mass divided by a twelfth of the mass of this isotope. ONLY for this isotope of carbon, is the mass number AND the atomic mass identical (12) So, why is the mass of other isotopes not exactly equal to mass number (# of protons and neutrons)? Answer: mass defect. Since it is the standard, 12C has a mass of 12u. BUT, if you were to assemble 28 nucleons (neutrons + protons) to make 28Si, the atom would have a mass less than 28 (27.97693u to be exact) Reason: atoms in a more energetic state (less stable) have some of their energy in the form of mass, and vice versa We don’t see this behaviour on the large scale (you don’t get lighter at a fast speed), but it happens at small scales! How do we know what isotopes a sample consists of? You can determine the isotopic abundance of a sample using a mass spectrometer (we don’t have one here but most universities would):  DETERMINING ATOMIC MASS OF ELEMENTS Imagine you are collecting donations for an event with 12 of your friends. 10 friends give you 20 dollars, but 2 of your friends are a bit short and only have 10 dollars. How would you determine the average donation? Average donation = (10 x $20 + 2 x $10) / 12 = $18.33 This is like the method for calculating atomic mass of an element (like those seen underneath elements on the periodic table) Equation: (2 ways to Relative atomic mass (Ar) = (% abundance of isotope 1 x mass of isotope 1) + (% abundance of isotope 2 x mass of isotope 2) + ….. calculate the 100 same thing) OR Ar = (decimal abundance of isotope 1 x mass of isotope 1) + (decimal abundance of isotope 2 x mass of isotope 2) + …  Why? +/+ repulsion overcomes the Some isotopes are unstable and tend to break apart strong nuclear force, or too through radiation after a certain amount of time… much energy in the nucleus NUCLEAR RADIATION Key Terms (distinguish) Radioactive Decay - the spontaneous disintegration of an unstable isotopes Nuclear radiation is energy or small particles emitted from a radioisotope as it decays. Radioisotopes is an isotope that spontaneously decays to produce 2 or more smaller nuclei and radiation. Radioactive means that a substance has the potential to emit nuclear radiation on decay. Periodic Trends Now that you’ve familiarized yourself a bit more closely with the periodic table, let’s look at some trends… …and try to give reasons for them based on # of protons in the nucleus and the location of the outermost electrons You’ll need to demonstrate this after the lesson through an activity. (New informational slide added after lesson) TWO factors impact the following trends we observe: 1. Electrostatic attraction between outermost electron and nucleus 2. “Shielding” effect of internal electrons - slight (-)/(-) repulsion which weakens the previously-mentioned attraction (This is known as Zeff: effective nuclear charge. This can be approximated through calculation for any electron - the overall electrostatic force holding the electron in, taking into account charge of nucleus and surrounding electrons that repel. However, we only look at these trends qualitatively this year)  Definition: Distance from nucleus to just beyond outermost electrons Atomic Radius Ionic Radius Isoelectronic series radius trend Why are these getting smaller? What is the same? What’s different?  Definition: Quantity of energy required to remove a (1st) single valence electron from an atom/ion in gaseous Ionization Energy state  Definition: The energy change that occurs when an atom in the gaseous state gains an electron Electron Affinity Same trend as last one, but the definition has a subtle difference (Preview - we’ll do a separate lesson on this) New Electronegativity term/category! Definition - How strongly an element will attract electrons in a bond What’s the significance of this? Depending on the EN difference, electrons will either be -Shared evenly -Shared but closer to one side -Completely taken to one side (Ionic) Metal-nonmetal relationships have a big electronegativity difference: electrons grabbed (ionic)  Called a “crystal lattice” structure Structure - Ionic Compounds  Chemical Formula is represented as a “FORMULA UNIT” Ionic compounds are represented by the ratio of ions… this is why we always simplify. Ex. MgO (1 - to - 1 ratio needed to balance charge) Compared with molecular compounds (next) which must show every atom in the molecule in the formula. ○ Ex. Vinegar (acetic acid) is CH3COOH ○ Ex. Ethane (dicarbon hexahydride) is C2H6 (not reduced) We don’t reduce with molecular - it’s necessary to represent EVERY atom in the structure with the formula Some key Ionic Compound Properties Takes MASSIVE amounts of energy to overcome this attraction High Melting and Boiling points Water is POLAR. It has + and - sides. Often Soluble in Water These break apart the lattice through attraction. Conduct Electricity when Dissolved Having ions separated in water allows them to move in order to keep current flowing Hard & Brittle Applying a force and shifting the lattice causes -/- and +/+ interactions that break it apart Naming Naming ionic compounds is EASY! International system of chemistry notation standards (International Union of Pure Simply: and Applied Chemistry) Name of metal + name of nonmetal but with -ide suffix NaCl: sodium chloride MgBr2: magnesium bromide …regardless of formula unit ratio! You don’t need prefixes, like dibromide, etc.  Multivalent Metals Many elements can exist in multiple different stable valence configurations. Multivalent metal ionic compounds need their charge specified with a roman numeral Ex: copper (I) chloride OR copper (II) chloride Formulas? Polyatomic ions Sometimes ions are multiple elements COVALENTLY (later) bonded together, which overall lack electrons or have extra in their molecular structure. Example: hydroxide (OH-) In a formula unit , if there are multiple of these, we MUST contain them in brackets Ca(OH)2 vs. NaOH Molecular Element (diatomic elements) H2 O2 These seven elements, when on their own, will covalently bond to themselves in order to complete the Br2 valence level. F2 I2 N2 Cl2 These could be single, double, or triple covalent bonds Molecular Element (diatomic elements) Significance: When we start chemical reactions (unit 2), you’ll need to complete reactions by expecting which products will form (based on the patterns we look at). If one of those elements (HOBrFINCl) is on its own, remember to write it with a 2 subscript! Covalent bonding Both atoms feel stable like a noble gas! FOLLOW ALONG ON PAPER TOGETHER (Let’s say you’re given CH2O) Steps for drawing a molecule *Note - this can be a challenge, but USUALLY we can expect how things will bond if you follow these steps* 1. Arrange the symbols so that the atom with the greatest number of unpaired electrons is in the center 2. At this point, you can sometimes expect what will bond where. The following steps present a systematic method: 3. Count all valence electrons 4. Place two electrons between the central atom and each of the surrounding atoms 5. Place lone pairs of electrons around each of the surrounding atoms to complete octet (except hydrogen) 6. Determine how many electrons are remaining from the original count 7. Place remaining electrons on the central atom in pairs 8. Check that the central atom has satisfied its octet. If not, move one of the lone pairs of the surrounding atoms into one of the existing bonding positions 9. Check that all octets are satisfied (2 for hydrogen) 10. Replace bonding pairs of electrons with a line between atoms to represent the covalent bonds FOLLOW ALONG ON PAPER (Let’s say you’re given NH4+) Steps for drawing a polyatomic ion *Note - this can be a challenge, but USUALLY we can expect how things will bond if you follow these steps* 1. Arrange the symbols so that the atom with the greatest number of unpaired electrons is in the center 2. At this point, you can sometimes expect what will bond where. The following steps present a systematic method: 3. Count all valence electrons and add/subtract from your total based on ion charge 4. Place two electrons between the central atom and each of the surrounding atoms 5. Place lone pairs of electrons around each of the surrounding atoms to complete octet (except hydrogen) 6. Determine how many electrons are remaining from the original count 7. Place remaining electrons on the central atom in pairs 8. Check that the central atom has satisfied its octet. If not, move one of the lone pairs of the surrounding atoms into one of the existing bonding positions 9. Check that all octets are satisfied (2 for hydrogen) 10. Replace bonding pairs of electrons with a line between atoms to represent the covalent bonds 11. Put square brackets around the structure and the ion charge in the top right Properties Attraction between molecules is relatively weak compared to the attraction between ions - therefore lower melting + boiling points Review the table on page 69 of the text which compares the properties of ionic and molecular compounds ○ Hint: they all relate to the weaker attraction between particles. This attraction still exists though and is important! More on this with intermolecular forces.  Check out Benzene - and example of an “aromatic” Exceptions to the Octet Rule compound. Some electrons are delocalized (not The octet rule is a simplification. confined to a single bond location) Some elements can form compounds by overfilling or underfilling valence shell - subsequent courses that look at quantum model structure will address this. nitrogen monoxide boron trichloride sulfur hexafluoride I will not give you one of these to draw on a test this year as I don’t etc… want to give you a headache.  Electronegativity The degree to which an individual atom, when bonded, will attract bonding electrons to itself Should fit logically with ionization energy/ e. affinity trend! “Pauling” scale, named after Linus Pauling - no unit  If the electrons are pulled to one side, it affects the overall distribution of +/- charge across the bond. Check these charge distribution colour maps: Varies depending on source Threshold for ionic is ~1.8-2.0 The bond character (ionic, polar covalent or nonpolar covalent) is determined by the difference in electronegativity numbers If the electronegativity difference is between 0 - 2.0 (ish), it’s a polar covalent bond. This leads to a bond dipole - think two poles, like North and South, but it’s (-) and (+) Just because you have a polar bond, it doesn’t mean the molecule is polar overall “So we’re polar, right?” “Just between you and me, yes… but we also need to think of the whole molecule” Water Although we draw lewis diagrams in 2D on the board , it’s important to realize that we live in a 3-dimensional world (I really need to get a volumetric display…) Q: Think back to the carbon atom in the molecular model kit - where were the bonding locations positioned? Why is it bent? (Impossible for the molecule to be linear with two single bonds) 2D representation Reality The C=O bond is polar, but the two bond dipoles cancel each other out since they are in opposite directions (linear) No “net dipole moment” - the two individual bond dipoles cancel out Water, on the other hand, with its bent shape, has a net dipole moment since both individual bond dipoles don’t entirely cancel out. DEMO - CAN I HAVE A VOLUNTEER WHO DOESN’T MIND MESSING THEIR HAIR UP? Why is this molecule non-polar? C H Q: you made ethane. What is The lowercase greek letter delta is used with +/- this molecule called? symbols as seen above to show sides of the dipole (hint: prefix for 6 is often hex) Task Methane vs. Methanol Which one is more polar, and why? In other words, which one is more likely to bend with Mr. Tillmann’s balloon demo? Chapter 3 - Molecules & Intermolecular Forces Review - before starting this lesson: What makes a molecule polar? Non-polar? Answer checklist: -Electronegativity and bond polarity -Overall molecule symmetry to determine molecule polarity Group discussion Can you explain why each of these molecules is P/NP? Remember the stream of water? Remember the stream of water? Because water is a polar molecule, the molecular dipoles rearrange and are attracted to a charged object Covalent bonding is an INTRAMOLECULAR force Intra = within/inside This bond is WITHIN the molecule of water  BUT! Water is a liquid at room temperature… Even though the particles are moving fast, something keeps them held together… The negative and positive ends of the dipoles are attracted… this is an example of an INTERMOLECULAR FORCE (+/- sided) POLAR MOLECULES HAVE DIPOLE-DIPOLE FORCES (Symmetrical charge) NONPOLAR MOLECULES HAVE DISPERSION FORCES Dispersion forces are more significant if molecules can pack closely together Example: sources of fats high in unsaturated fatty acids are typically liquids at room temperature - the molecules can’t pack together as closely, therefore less opportunity for dispersion forces to hold them together Saturated = every carbon is saturated with hydrogen. Unsaturated means a double bond exists instead of hydrogens  CONTINUE HERE Molecules with dipoles due to the following bonds: H-O H-F H-N H-(atom with high EN) Are given a special name for their dipole-dipole interaction since it’s particularly strong: Hydrogen Bond Theory Water is attracted to polar substances such as glass through hydrogen bonding. But if water were to interact with a nonpolar substance, such as wax, it won’t adhere. This is why water beads up and slips off a freshly waxed car. (Freshly waxed car surface)  Do you notice this behaviour at the edge of narrow glassware when you fill it with water? This is the beginning of capillary action. The water molecules near the edge are attracted to the dipoles of glass molecules (SiO2) Watch for my capillary tube demo during the competition