Download presentation

Presentation is loading. Please wait.

Published byOphelia Bridges Modified over 6 years ago

1
Atomic Structure

2
Subatomic Particles (Particles that make up an atom) ● Proton (p+) - Positively charged - Found in the nucleus - Large mass ●Neutron (n 0 ) ● Neutron (n 0 ) - A neutral particle - Found in the nucleus - Large mass ●Electron (e-) ● Electron (e-) - Negatively charged particle - Found outside of the nucleus in the electron cloud - Very small mass

3
Summary ●The nucleus has almost all of the mass & it has a + charge ● The nucleus has almost all of the mass & it has a + charge ●The electron cloud has a – charge & creates the atom’s ● The electron cloud has a – charge & creates the atom’s volume volume

4
How to read a box on the periodic table 11 Na 22.98 ● Atomic Number - # above symbol - Always determines # of protons (can never change!) - Determines # of electrons if atom is neutral (0 charge) - We assume the periodic table is neutral (same # of p+ & e-) ● Summary: - Sodium’s atomic number is 11 - Sodium has 11 protons & 11 electrons Atomic # Symbol

5
11 Na 22.98 ● Average atomic mass - # below the symbol in decimal form - The average mass of an atom - Not all sodiums have the same mass due to different number of neutrons (isotope) Average atomic mass

6
● Mass Number – rounding the a.a.m. to a whole number - Mass # = # of protons + number of neutrons - Therefore, use to find number of neutrons mass # - # of p = # of n ● Summary - Na’s ave. atomic mass = 22.98 amu (atomic mass units) - Na’s mass # = 23 - Number of neutrons in Na: 23 – 11 = 12 neutrons 11 Na 22.98 Mass # (23)

7
You just have to try one! 47 Ag 107.87 Determine: 1. Atomic # = 2. # of protons = 3. # of electrons = 4. Ave. atomic mass = 5. Mass number = 6. # of neutrons =

8
Isotopes ● Atoms of the same element can have different numbers of neutrons, therefore, different masses - Remember, neutrons have mass! - Changing the number of neutrons, changes the mass ● Let’s look at 2 isotopes of carbon as an example: - Carbon ALWAYS has 6 protons - But it can have a mass of 12 amu (6p + 6n) C or C-12 - and it can have a mass of 14 amu (6p + 8n) C or C-14 12 14 6 6

9
Perfect practice makes perfect! ● Here is an isotope of oxygen: O - How many protons are present? __________ - What is the mass number? __________ - How many neutrons are present? __________ - How many electrons are present? __________ ● Write the shorthand form of a nitrogen isotope that has 13 neutrons. _ N or N - __ 18 8 _

10
Mole Conversions ● Moles (mol) are a unit of measurement ● 1 mole = 6.02 x 10 23 units (atoms, molecules, formula units, ions, etc) ● 6.02 x 10 23 is Avogadro’s number ● Mole Conversions 1 mole = 6.02 x 10 23 units = formula weight (grams)

11
What is formula weight? ● Formula weight is the weight of an element or compound in grams ● How is formula weight determined? - Use your periodic table and find the ave. atomic mass - Formula weight of H 2 O - H’s ave. atomic mass = 1.01 g (x 2) = 2.02 g - O’s ave. atomic mass = 16.00 g 18.02 g H 2 O 2.02 g + 16.00 g = 18.02 g H 2 O

12
What is the formula weight of… Al Al Br 2 Br 2 MgF 2 MgF 2 CH 4 CH 4 Ca 3 (AsO 4 ) 2 Ca 3 (AsO 4 ) 2

13
Conversions 1. Moles to grams # of moles x formula weight (g) = _____ grams 1 1 mole ● Example: How many grams are in 5.00 moles of CaCl 2 ? Formula weight of CaCl 2 : ● Ca = 40.08 g Cl = 35.45 g (x2) = 70.90 g ● 40.08 g + 70.90 g = 110.98 g CaCl 2 5.00 moles x 110.98 g CaCl 2 = 554.9 = 1 1 mole 555 g CaCl 2

14
2.Grams to moles # of grams x ___1 mole _ = _______ moles 1 formula wt (g) ● Example: How many moles are in 25.00 g of NaCl? 25.00 g of NaCl x _ 1 mole___ = 1 58.44 g NaCl 0.4278 moles of NaCl

15
3.Moles to units (atoms, molecules, formula units, ions, etc.) # of moles x 6.02 x 10 23 units = ____ units 1 1 mole ● Example: How many atoms are in 0.250 moles of neon? 0.250 moles of Ne x 6.02 x 10 23 atoms = 1 1 mole 1.51 x 10 23 atoms of Ne

16
4. Units to moles # of units x ___1 mole____ = ____ moles 1 6.02 x 10 23 units ● Example: How many moles are in 4.23 x 10 24 molecules of H 2 O? 4.23 x 10 24 molecules x ______1 mole______ = 1 6.02 x 10 23 molecules 7.03 moles of H 2 O

17
5. Grams to units # of grams x 6.02 x 10 23 units = ____ units 1 formula wt (g) ● Example: How many formula units are in 35.0 g of K 2 O? 35.0 g K 2 O x 6.02 x 10 23 formula units = 1 94.20 g K 2 O 2.24 x 10 23 formula units of K 2 O

18
6. Units to grams # of units x _formula wt (g)_ = ____ grams 1 6.02 x 10 23 units ● Example: How many grams are in 9.75 x 10 25 atoms of Ag? 9.75 x 10 25 atoms x __107.87 g Ag__ = 16.02 x 10 23 atoms 17500 g Ag

19
Mass Percent Composition ● Determining what percentage of each element is in a specific formula ● Example: Find the mass % of each element in NaHCO 3. - Step 1: Find their individual ave. atomic masses from the PT & multiply by the number of atoms of each (subscript) Na = 22.99 g (1)= 22.99 g H = 1.01 g (1) = 1.01 g C = 12.01 g (1) = 12.01 g O = 16.00 g (3) = 48.00 g 84.01 g NaHCO 3 - Step 2: Add them to get the total weight of the formula.

20
- Step 3: Find the mass % of each! -Remember:Na = 22.99 g (1)= 22.99 g H = 1.01 g (1) = 1.01 g C = 12.01 g (1) = 12.01 g O = 16.00 g (3) = 48.00 g 84.01 g of NaHCO 3 - Take the elements individual total weight and divide by the total weight of the formula. Then Multiply by 100. - Mass % of Na = 22.99g /84.01 (100) = 27.36 % - Mass % of H = 1.01g /84.01 (100) = 1.20 % - Mass % of C = 12.01g /84.01 (100) = 14.30 % - Mass % of O = 48.00g /84.01 (100) = 57.14 % - Add %’s to make sure they add up to 100%

21
Getting the formula from mass % ● Do the opposite of finding the mass % ● Example: What is the formula of a substance that is made of 27.29% C & 72.71% O. The total weight of the substance is 44.01 g. - Step 1: Divide each % by 100 then multiply by the total weight C : 27.29/100 = 0.2729 (44.01 g) = 12.01 g C O: 72.71/100 = 0.7271 (44.01 g) = 32.00 g O - Step 2: Divide the totals by their average atomic mass (from PT) 12.01 g C/12.01 g C = 1 32.00 g O/16.00 g O = 2 - Step 3: Put the formula together CO 2

22
Finding the relative atomic mass ● Where does the periodic table get its average atomic masses from? ● Here’s an example: There are two isotopes of chlorine which consists of atoms of relative isotopic masses 35.0 (75.0 %) and 37.0 (25.0 %). % abundance Isotope mass Cl-3575.0 35.0 amu Cl-3725.0 37.0 amu (75.0/100) x 35.0 amu + (25.0/100) x 37.0 amu = 35.5 amu The answer matches Cl on the periodic table!

Similar presentations

© 2021 SlidePlayer.com Inc.

All rights reserved.

To make this website work, we log user data and share it with processors. To use this website, you must agree to our Privacy Policy, including cookie policy.

Ads by Google