Unit 1: Atomic Structure and Properties.

1.1 Moles and Molar Mass.

Convert between number of particles (atoms, molecules, formula units, ions) and moles, using Avogadro’s number.

  1. How many moles are in 3.01x1023 molecules of CO2?

Use molar mass to convert between mass and moles of a substance.

  1. How many moles are in 25.0 g of H2O?

Use subscripts to convert between molecules and the number of atoms of a given element.

  1. How many hydrogen atoms are in 2 molecules of NH3?

1.2 Mass Spectra of Elements.

Identify an element given its mass spectrum.

  1. What element is represented by the mass spectra?
    2. Mass spectrum example

Calculate the average atomic mass of an element given its mass spectrum.

  1. What is the average atomic mass of the element depicted in the mass spectrum below?
    Mass spectrum of Neon

Determine the number of protons, neutrons, and electrons in a given isotope.

  1. How many protons, neutrons, and electrons are in a neutral isotope of Carbon-12?

1.3 Elemental Composition of Pure Substances.

Determine the percent composition by mass of a compound.

  1. What is the percent composition by mass of each element in lactic acid (C3H6O3)?

Determine the empirical formula of a compound.

  1. What is the empirical formula of lactic acid (C3H6O3)?

Determine the empirical formula of a compound from percent composition by mass data.

  1. What is the empirical formula of a compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass?

Determine the formula of a compound from gravimetric analysis lab data.

  1. 2.0g of an unknown sulphate salt (XSO4) is reacted with excess BaCl2(aq). Barium sulphate precipitate (BaSO4(s)) was produced. The precipitate was carefully filtered, rinsed, dried, and its mass was determined to be 2.54g. Determine the identity of the unknown sulphate salt.
  2. Copper metal (Cu(s)) reacts with oxygen gas (O2(g)) to form a copper oxide when heated. Use the lab data to determine the chemical formula of the copper oxide.
      Mass of sample before heating : 5.00g
      Mass of sample after heating: 6.25g.

Determine the empirical formula of a hydrate from lab data obtained by heating the sample.

  1. Determine the empirical formula of the hydrate MgSO4 * xH2O from the lab data:
    Mass of empty crucible: 20.00g
    Mass of hydrate and crucible before heating: 24.90g
    Mass of sample after 1st heating: 24.35g
    Mass of sample after 2nd heating 24.11g
    Mass of sample after 2nd heating 24.10g

1.4 Composition of Mixtures.

Differentiate between a pure substance and a mixtures.

  1. Classify the following as a pure substance or mixture: (a) salt water (b) gold (c) air.

Convert between mass or the percent composition by mass of an element and the mass or the percent composition by mass of a compound in a mixture.

  1. A mixture of magnesium oxide (MgO) and sodium oxide (Na2O) contains the 16.3g MgO and 13.7g Na2O. What percent by mass of the mixture is magenesium (Mg)?
  2. A 100.0g mixture of magnesium oxide (MgO) and sodium oxide (Na2O) contains the 25.0% by mass sodium (Na). How many grams of MgO and Na2O are contained in the sample?

Identify a possible impurity in a sample based on the percent composition by mass of the sample and the pure compound.

  1. A students believes a sample of KCl maybe be comtaminated with NaCl or KF. They determine the percent composition by mass of K to be 50.0% Justify if the sample is impure what what the impurity is.

1.5 Atomic Structure and Electron Configuration.

Identify the location of subatomic particles and the charge of areas of the atom.

  1. Summarize the location, charge, and relative mass of the subatomic particles.

Understand and apply the terms: shell (or energy level), subshell (or sublevel), and orbital in reference to the quantum mechanical model.

  1. Describe the atomic structure of Neon (Ne) using the terms: shell, subshell, and orbital.

Qualitatively apply Coulomb’s law and effective nuclear charge to justify differences in attractive force between nucleus and core vs. valence electrons.

  1. Why are valence electrons easier to remove than core electrons?

Write the electron configuration for an atom or ion.

  1. Write the electron configuration for Cl
  2. Write the electron configuration for O2-.
  3. Write the electron configuration for Mn2+

Determine the number of unpaired electrons in an atom.

  1. How many unpaired electrons are in a nitrogen atom?

Differentiate electronic structure in terms of shells and subshells.

  1. What is the difference between the 2s and 2p subshells in terms of energy and shape?

Qualitatively apply Coulomb’s law to discuss differences in ionization energy by comparing electron shells/subshells between atoms.

  1. Why does lithium (Li) have a higher ionization energy than potassium (K)?
  2. Why does lithium (Li) have a lower ionization energy than nitrogen (N)?

1.6 Photoelectron Spectroscopy.

Identify an element based on its PES (Photoelectron Spectroscopy).

  1. Identify the element shown in the PES below

    2. PES of a unkown element

Label the peaks on a PES spectrum with the corresponding subshell.

  1. In a PES spectrum of oxygen below , label which peaks correspond to 1s, 2s, and 2p electrons. PES of a unkown element 2p,2s,1s

    2. PES of a unkown element

Compare heights of peaks in the spectrum to determine the number of electrons in a subshell.

  1. If the 2p peak is three times the height of the 2s peak in a PES spectrum, how many electrons are in each subshell?

Explain differences between PES spectra of two atoms/ions.

  1. How would the PES spectrum of Na differ from that of Na+?

1.7 Periodic Trends.

Apply Coulomb’s Law and atomic structure to compare: atomic radius, electronegativity, ionization energy, and electron affinity.

  1. Explain why fluorine (F) has a smaller atomic radius than oxygen (O).
  2. Predict whether sodium (Na) or lithium (Li) has a larger atomic radius.
  3. Explain why the first ionization energy of francium (Fr) is lesser than that of caesium (Cs)
  4. Predict if the first ionization energy of copper (Cu) is greater or lesser than that of iron (Fe)
  5. Explain why florine (F) has a greater electronegativity or iodine (I)
  6. Compare the electron affinity of oxygen (O) and carbon (C)

Compare the atomic radius a neutral atom to its cation or anion using Coulomb’s Law and atomic structure.

  1. Justify whether Na or Na+ has a larger atomic a radius

Compare the atomic radius of isolectronic (equal number of electrons) species using Coulomb’s Law and atomic structure.

  1. Rank the size of the atomic radii of Kr, Br -, and Rb+ from least to greatest.

1.8 Valence Electrons and Ionic Compounds.

Predict the type of bond that will form between two atoms based on their valence electrons.

  1. What type of bond is expected between magnesium (Mg) and chlorine (Cl)?
  2. What type of bond is expected between carbon (C) and chlorine (Cl)?

Apply the fact that Elements in the same group have analagous (similar) properties to solve problems.

  1. Which compound has properties most similar to NaCl: KCl, CaCl2, or ClF?

Predict the charge on an atom by its location on the periodic table.

  1. What the charge of the calcium ion?
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