1.6 Relative atomic mass

• The atomic mass refers to the mass of a single atom of an element
• Isotopes are different forms of the same element with the same number of protons but different numbers of neutrons
• Relative atomic mass eliminates the need to know the exact mass of each atom.


  True
• The relative atomic mass uses the atomic mass unit (amu) as a reference.
• The base unit of mass in the SI system is the kilogram
• Relative atomic mass is the average mass of atoms of an element compared to 1/12 the mass of a carbon-12 atom.
• What are the two key characteristics of relative atomic mass?

  Average mass and dimensionless
• What is another name for the atomic mass unit (amu)?

  Dalton
• Steps to calculate relative atomic mass using abundances in percentages

  1️⃣ Multiply the abundance by the mass for each isotope

  2️⃣ Sum the results

  3️⃣ Divide by 100
• The relative atomic mass (Ar) is calculated by summing the product of the abundance and the mass
• The relative atomic mass (Ar) is compared to 1/12 the mass of a carbon-12 atom.
• Relative atomic mass is the average mass of an atom compared to 1/12 the mass of a carbon-12 atom.
• What is the product of 75.8 × 34.969 in the relative atomic mass calculation for X-35?

  2649.94
• What is the relative atomic mass defined as?

  Average mass compared to carbon-12
• What is atomic weight defined as?

  Weighted average of isotopes
• How do isotopes affect relative atomic mass?

  Weighted average of masses
• Match the mass unit with its definition:

  Atomic Mass Unit ↔️ 1/12 the mass of carbon-12

  Kilogram ↔️ SI base unit of mass

  Gram ↔️ 1/1000 of a kilogram
• What formula is used to calculate relative atomic mass?

  Weighted average of isotopes
• How many grams are in a kilogram?

  1000
• What is the formula to calculate relative atomic mass (Ar)?

  $Ar =$$\sum_{i} (abundance_{i} \times mass_{i})$
• Relative atomic mass allows us to compare the masses of different elements on a common scale.
• Match the isotope with its properties:

  Carbon-12 ↔️ 6 protons, 6 neutrons

  Carbon-14 ↔️ 6 protons, 8 neutrons
• The relative atomic mass uses the amu as a reference.
• The abundance of isotopes is given as percentages in relative atomic mass calculations.


  True
• Relative atomic mass is a dimensionless quantity with no units.


  True
• Relative atomic mass takes into account the masses and natural abundances of all isotopes.

  True
• After summing the values in a relative atomic mass calculation, the result is divided by 100 because the abundances are given as percentages
• The relative atomic mass is a dimensionless quantity with no units.


  True
• Relative atomic mass allows us to compare the masses of different elements on a common scale.


  True
• The relative atomic mass is compared to 1/12 the mass of a carbon-12 atom.
• What is the atomic mass unit (amu) defined as?

  1/12 the mass of carbon-12
• Relative atomic mass compares the masses of elements to carbon-12.


  True
• Steps to calculate relative atomic mass

  1️⃣ Identify isotopes and their abundances

  2️⃣ Multiply abundance by mass for each isotope

  3️⃣ Sum up the values

  4️⃣ Divide by 100 if abundance is in percentages
• What are isotopes of an element?

  Same protons, different neutrons
• Isotopes affect relative atomic mass by averaging their masses according to their natural abundance.


  True
• What is the formula to calculate relative atomic mass (Ar)?

  Ar = \sum_{i} (abundance_i \times mass_i)</latex>
• What formula is used to calculate the relative atomic mass (Ar) of an element?

  $Ar =$$(abundance × mass) +$$...$
• Why is the sum of the (abundance × mass) values divided by 100 in relative atomic mass calculations?

  To convert to percentages
• How does relative atomic mass help in stoichiometric calculations?

  Compares element masses
• Steps to calculate relative atomic mass (Ar)

  1️⃣ Multiply the abundance by the mass for each isotope

  2️⃣ Sum up the values

  3️⃣ Divide by 100 if abundances are percentages