QuÃ_mica Inorgánica UMG 2025
Page 1: Content Information
Document Details
Title: QUÍMICA INORGÁNICA
Author: Lic. Mario Santiago Zetina
Year: 2012
Page 2: Basic Equations
Key Formulas
Media Velocity Formula:
v = d/t
Rearranging for t:
t = d/v
Density Formula:
ρ = m/V
Rearranging for m:
m = ρ \cdot V
Ohm's Law:
V = I \cdot R
Rearranging for I:
I = V/R
Kinetic Energy Formula:
E_k = 1/2 \cdot mv²
Rearranging for v:
v = √(2E_k/m)
Page 3: More Equations
Additional Formulas
Force Formula:
F = m \cdot a
Rearranging for a:
a = F/m
Circle Area Formula:
A = πr²
Rearranging for r:
r = √(A/π)
Specific Heat Capacity:
Q = m \cdot c \cdot ΔT
Rearranging for c:
c = Q/(m \cdot ΔT)
Universal Gravity Equation:
F = G(m₁m₂/r²)
Rearranging for r:
r = √(G(m₁m₂)/F)
Page 4: References
Textbook Reference:
Title: QUÍMICA
Edition: 13ª edición
Authors: Raymond Chang and Jason Overby
Publisher: McGraw Hill
Page 5: History of Chemistry
Historical Milestones
Prehistory: Discovery of fire, pottery, and paint production.
Antiquity: Greek philosophers adopt the concept of matter.
Alchemy: Search for the philosopher's stone and the elixir of life.
Page 6: Key Figures and Developments
Significant Contributions
Iatrochemistry: Medicinal chemistry.
Antoine Lavoisier (1776): Known as the "father of chemistry"; detailed explanation of combustion and scientific method development.
Modern Chemistry: Discovery of chemical elements, atomic theory, organic chemistry, etc.
Page 7: What is Chemistry?
Definition and Scope
Chemistry is the study of matter and its changes.
It deals with substances and their interactions.
Focuses on preparation, properties, and transformations of substances.
Page 8: Branches of Chemistry
Major Branches
Pure Chemistry: Studies all types of substances, organic and inorganic.
Organic Chemistry: Focused on carbon compounds.
Inorganic Chemistry: Studies structure, composition, and reactions of elements and compounds.
Analytical Chemistry: Analyzes properties of chemical systems.
Industrial Chemistry: Examines production methods of chemical substances.
Page 9: Basic Concepts
Definitions
Matter: Anything that occupies space and has mass (visible and invisible).
Substance: A form of matter with a defined composition and distinct properties (e.g., water, sugar).
Mixture: Combination of two or more substances retaining their properties (e.g., milk, cement).
Page 10: Types of Mixtures
Classification of Mixtures
Homogeneous Mixtures: Uniform composition (elements indistinguishable).
Heterogeneous Mixtures: Non-uniform composition (distinguishable substances).
Page 11: Examples of Homogeneous Mixtures
Visual Representations
Example: Saltwater is a homogeneous mixture.
Visual Representation: Figure illustrating uniformity of saltwater solution.
Page 12: Examples of Heterogeneous Mixtures
Visual Representations
Visual Representation: Image illustrating distinct phases of heterogeneous mixtures.
Page 13: Identifying Mixtures
Examples and Questions
Gold and silver alloy (homogeneous).
Activated carbon in a liquid (heterogeneous).
Carbon dioxide in water (homogeneous).
Silica and clay in water (heterogeneous).
Soil with small stones (heterogeneous).
Page 14: Separation of Mixtures
Key Concept
Mixtures can be formed and separated through physical means without changing the identity of components (e.g., boiling saltwater).
Page 15: Elements vs. Compounds
Definitions
Element: A substance that cannot be separated into simpler substances by chemical means.
Compound: A substance formed from two or more elements chemically bonded in fixed proportions.
Separation: Compounds only separable by chemical methods.
Page 16: Common Elements and Symbols
Element List
Aluminum (Al), Arsenic (As), Sulfur (S), Barium (Ba), Bismuth (Bi), Bromine (Br).
Page 17: Diagrams of Elements and Compounds
Representation
Visual representations of atoms indicating their classification as elements or compounds.
Page 18: Classification of Matter
Hierarchical Structure
Matter can be separated into:
Mixtures
Homogeneous
Heterogeneous
Substances
Compounds
Elements
Page 19: States of Matter
Three States
Solid
Liquid
Gas
Page 20: Characteristics of States of Matter
Properties
Solid: Atoms/molecules are closely packed, minimal movement.
Liquid: Atoms/molecules close but not rigid; can move.
Gas: Atoms/molecules are spread apart; large distances between them.
Page 21: Questions about States of Matter
Inquiry Questions
Why are gases more compressible than liquids or solids?
What happens molecularly when solids change to liquids (melting)?
Why do solids have defined shapes while liquids and gases do not?
Page 22: Phase Changes
Key Points
Melting Point: Temperature where solid becomes liquid.
Boiling Point: Temperature where liquid becomes gas.
Condensation Point: Temperature where liquid becomes solid.
Page 23: Physical and Chemical Properties
Types of Properties
Physical Property: Can be observed/measured without changing substance (color, melting point).
Chemical Property: Alters molecular structure during interactions (e.g., hydrogen combustion).
Page 24: Chemical Change
Concept Overview
Reaction that leads to the formation of products with different properties from the reactants.
Page 25: Examples of Properties
Property Identification
Iron forming rust (chemical).
Metal stretching into wire (physical).
Temperature at which liquid becomes gas (physical).
Copper reacts with sulfuric acid (chemical).
Vinegar and baking soda reaction (chemical).
Wood burning (chemical).
Reflective surface (physical).
Cooling reduces a substance's volume (physical).
Page 26: Extensive vs. Intensive Properties
Definitions
Extensive Properties: Depends on amount of substance (e.g., mass, volume).
Intensive Properties: Independent of amount of substance (e.g., temperature, density).
Page 27: Properties Comparison
Summary of Properties
Intensive: Temperature, boiling point, density.
Extensive: Weight, size, volume.
Page 28: Diagrams of Elements
Visual Comparison
Diagrams portraying compounds vs. physical/chemical changes.
Page 29: Notation
Data Presentation
Notation information placeholder (context needed).
Page 30: Measurements
Types of Properties
Macroscopic Properties: Can be directly measured.
Microscopic Properties: Cannot be directly observed.
SI Units
Time
Mass
Volume
Density
Temperature
Page 31: Measurement Tools
Common Instruments
Measuring flask, graduated cylinder, burette, pipette.
Page 32: SI Units Overview
Basic Units
Length: Meter (m)
Mass: Kilogram (kg)
Time: Second (s)
Electric Current: Ampere (A)
Temperature: Kelvin (K)
Amount of Substance: Mole (mol)
Luminous Intensity: Candela (cd)
Page 33: Mass vs. Weight
Key Differences
Mass: Amount of matter in an object.
Weight: Gravity's force on an object.
SI Units
Mass: kg, Volume: m³.
Page 34: Density Exercises
Practical Exercise Example
Cylinder A: 800 g, 1000 cm³, Density = 0.8 g/cm³
Cylinder B: 1000 g, 2000 cm³, Density = 0.5 g/cm³
Conclusion: Cylinder A has higher density.
Page 35: Gold Properties
Example Calculation
Gold sample: 301 g, Volume: 15.6 cm³
Density = 19.3 g/cm³.
Page 36: Density Practice Problems
Exercises
Platinum density problem.
Mercury density problem.
Page 37: Temperature Scales
Key Temperature Points
Fahrenheit: Freezing 32°F, Boiling 212°F.
Celsius: Freezing 0°C, Boiling 100°C.
Kelvin: Absolute temperature scale.
Page 38: Temperature Conversion Exercises
Sample Conversion Problems
Convert Celsius to Fahrenheit (80 °C).
Convert Fahrenheit to Celsius (130 °F).
Page 39: Scientific Notation
Overview
Scientific notation representation details for calculations.
Page 40: Scientific Notation Application
Importance
Used to simplify multiplication/division of large/small numbers in chemistry.
Page 41: Finding N in Scientific Notation
Explanation
Description of moving decimal for scientific notation.
Page 42: Special Cases in Notation
Concept Overview
Handle special cases in scientific notation.
n = 0 implication.
Page 43: Arithmetic with Scientific Notation
Examples of Operations
Addition, subtraction, multiplication of scientific notation formats.
Page 44: Significant Figures
Representation of Figures
Different notations for significant figures
Page 45: Definition of Significant Figures
Explanation
The last digit in a measurement is always uncertain, indicating precision of measurements.
Page 46: Rules for Significant Figures
Specific Guidelines
Non-zero digits are significant.
Captive zeros are significant.
Leading zeros are not significant.
Trailing zeros are significant if there's a decimal present.
Page 47: Significant Figure Problems
Practice
Determine significant figures given various numbers.
Page 48: Rules for Calculations
Operations with Significant Figures
Addition/Subtraction: Depends on decimal places.
Multiplication/Division: Depends on significant figures.
Page 49: Additional Calculation Guidelines
Redundant Rules
Specifics on how to handle rounding in significant figures during addition and subtraction.
Page 50: Accuracy vs. Precision
Definitions
Accuracy: Closeness to true value.
Precision: Closeness among multiple measurements.
Page 51: Example Scenario
Measurement Task
Comparing students’ mass measurements for accuracy and precision.
Page 52: Precision and Accuracy Test
Example Problem
Measurement task showcasing accuracy and precision in science.
Page 53: Conceptual Test
Understanding Inquiry
Evaluating performance in terms of accuracy or precision.
Page 54: Measurement Data Comparison
Data Presentation
Players’ scores representing precision and accuracy comparisons.
Page 55: Measuring Metal Mass
Measurement Results
Series of mass measurements for a metal cube and evaluation of accuracy and precision.
Page 56: Conversion Factors
Key Concept
Conversion factors facilitate transitioning between measurement units.
Page 57: Practical Application of Conversion Factors
Example Exercises
How many cm in 5.84 m?
Convert kg to lb.
Glucose daily intake conversion.
Volume conversion for helium.
Density of mercury in kg/m³.
Page 58: Additional Conversion Exercises
Sample Problems
Convert gallons to cubic centimeters.
Density to volume.
Page 59: Final Conversion Exercises
Measurement Practices Example
Convert gallons to cubic centimeters.
Temperature and density calculations.
Page 60: Thermochemistry
Fundamental Concepts
Energy capacity in work production.
Page 61: Types of Energy
Energy Classifications
Kinetic, radiant, thermal, chemical, and potential energy.
Page 62: Energy in Reactions
Energy Transfer
All chemical reactions absorb or release energy, often as heat.
Page 63: Additional Concepts
Placeholder Content
Page 64: Thermodynamic Systems
Systems Classifications
Open, closed, and isolated systems defined by energy and mass transfer.
Page 65: Exothermic vs. Endothermic Reactions
Energy Changes
Definitions and examples of energy release and absorption in reactions.
Page 66: Application of Energy Concepts
Understanding Systems
Classifying systems and reactions based on energy characteristics.
Page 67: System Evaluation
Problem-Solving
Classify system types and determine reaction types for given scenarios.
Page 68: First Law of Thermodynamics
Core Principles
Energy cannot be created or destroyed, only transformed.
Page 69: Calorimetry
Overview of Measurements
Calorimetry, specific heat, and capacity heat definitions and formulas explained.
Page 70: Specific Heat Calculation Example
Sample Calculation
Calculation of heat based on specific heat values.
Page 71: Heat Transfer Calculations
Equations for Heat Calculation
Detailed equations describing heat absorption/release.
Page 72: Heat Absorption Example
Practical Calculation Problem
Example calculating heat absorbed by water during temperature increase.
Page 73: Copper Heat Calculation
Heat Calculation Scenario
Example calculating heat necessary to change the temperature of copper.
Page 74: Heat Release Calculation Problem
Example Heat Calculation
Evaluating heat released during a cooling process.
Page 75: Mass Calculation Example
Assessing Heat Capacity
Problem determining mass based on specific heat and temperature change.
Page 76: Initial Temperature Determination
Problem-Solving Guide
Task determining the initial temperature of a heated aluminum sheet.
Page 77: Short Exam Example
Test Questions
Set of exercises demonstrating measurement and conversion skills in chemistry.
Page 78: Short Exam Instructions
Exam Content
Guidelines and examples for answering chemistry measurement problems.