Physics in Anaesthesia
- ISBN-10:
- 1904842984
- ISBN-13:
- 9781904842989
- Pub. Date:
- 04/01/2012
- Publisher:
- Scion Publishing Ltd.
- ISBN-10:
- 1904842984
- ISBN-13:
- 9781904842989
- Pub. Date:
- 04/01/2012
- Publisher:
- Scion Publishing Ltd.
Physics in Anaesthesia
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Overview
Physics in Anaesthesia covers the core physics knowledge required by anyone working in anesthesiology. It assumes only limited prior knowledge and has been written especially for those who consider themselves non-physicists. It covers material in an informative and accessible way from the very basics, and provides an important link between theory and practice.
Worked examples highlight the relevance to clinical practice, and along with graphs and charts, make the basics of physics understandable not only to doctors, but also to anesthesiologist assistants, nurse anesthetists and students.
Two different types of self-assessment questions (single best answers and multiple choice) at the end of each chapter test the reader’s understanding of the key concepts, while a summary section for each topic is ideal as a rapid refresher.
Product Details
| ISBN-13: | 9781904842989 |
|---|---|
| Publisher: | Scion Publishing Ltd. |
| Publication date: | 04/01/2012 |
| Edition description: | 1 |
| Pages: | 384 |
| Product dimensions: | 6.70(w) x 9.50(h) x 0.90(d) |
Read an Excerpt
Physics in Anaesthesia
By Ben Middleton, Justin Phillips, Rik Thomas, Simon Stacey
Scion Publishing Limited
Copyright © 2012 Scion Publishing LtdAll rights reserved.
ISBN: 978-1-907904-11-0
CHAPTER 1
Atoms and matter
Having read this chapter you will be able to:
• Appreciate the planetary and Bohr models of the atom.
• Define an element's atomic number and atomic mass.
• Recall the key differences between solids, liquids and gases.
• Understand the role of energy in changing states.
• Recognize the value of phase diagrams for showing state, triple point and critical point.
1.1 The atom
The word 'atom' originates from the Greek atomos meaning indivisible. In 1912, however, a New Zealand physicist, Ernest Rutherford, caused a sensation by revealing the atom is divisible. He had shown that the mass of the atom is concentrated in a tiny positively charged nucleus, surrounded by a tenuous cloud of negatively charged electrons. The new science of atomic physics was born which, for better or worse, heralded the beginning of the atomic age.
Rutherford's model of the atom
Rutherford created what is now the classic model of the atom, that of an 'atomic planetary model' with the electrons (planets) orbiting the nucleus (the sun). Planets are drawn to the sun due to gravitational force, but the attraction for the atom is due to the particles' electrical charges; the positively charged nucleus attracts the negatively charged electrons.
The nucleus of the atom contains nucleons, and is where virtually all of the atom's mass is held. There are two types of nucleons: protons and neutrons, and both have approximately the same mass, which is about 1840 times the mass of an electron. Protons are positively charged while neutrons have no charge. The number of protons defines the atomic number and may be thought of as the 'fingerprint' of an element because it is fixed for a specific element, e.g. hydrogen has atomic number = 1 and carbon has atomic number = 6. The atomic number determines the element's place in the periodic table.
The total number of nucleons is almost exactly equal to the atomic mass. Atomic mass is expressed in units of atomic mass units (not in units of actual mass). There are small differences between the atomic mass and the nucleon number depending on the element in question.
Atoms of the same element can have different numbers of neutrons in their nucleus, and these are known as isotopes of the element. For example, helium (atomic number = 2) has two isotopes: helium-3, and helium-4. Helium-4 is by far the most common isotope and has two protons and two neutrons in its nucleus (see Figure 1.1), so has an atomic mass of approximately 4. Helium-3, which is highly sought after for fusion research, has only one neutron so has an atomic mass of approximately 3. On earth, there are less than two atoms of helium-3 for every 10 000 of helium-4.
Some isotopes are stable while some are highly unstable and emit particles or radiation as they disintegrate. These isotopes are described as radioactive and are discussed in more detail in Chapter 26.
Units. Atomic mass number was originally standardized so that one atomic mass unit was equal to the mass of a proton (a hydrogen nucleus). This convention has now been changed so that an atomic mass unit is equal to 1/12th of the mass of a carbon-12 nucleus.
The Bohr model and energy levels
Just two years after publication of Rutherford's model of the atom, the Danish physicist Niels Bohr incorporated the idea of energy levels into a new atomic model. Bohr's model had strict rules for electrons: they could only exist in defined energy levels, so they could jump from one level to another but their energy levels were fixed. These energy levels are organized into 'shells' around the atom, an idea which forms a cornerstone of quantum physics and led to the development of the laser (see Chapter 24). Sometimes Bohr's model is called the Rutherford–Bohr model as Bohr essentially improved Rutherford's original model.
Chemical bonding
The attraction between atoms is known as a chemical bond, which allows the formation of chemical substances containing two or more atoms. Chemical bonds can be strong interatomic bonds such as covalent bonds or ionic bonds, or (usually) weaker intermolecular bonds such as dipole–dipole interactions or hydrogen bonding (see Section 2.4).
In covalent bonds, two atoms share one or more of their outer shell electrons. The negatively charged electrons occupy the space between the positively charged nuclei and are attracted to both nuclei simultaneously. The electrons can be thought to exist in a 'cloud' between the nuclei, because they are moving rapidly around an equilibrium position between the atoms. This attraction overcomes the repulsion which would otherwise exist between the two nuclei, so a strong bond is formed. Covalent bonds usually form between non-metallic atoms, for example, in organic compounds, as well as in diatomic gases and water molecules.
In an ionic bond, an outer electron is transferred from one atom to another. The electron is more tightly bound in the new atom so is able to exist at a lower energy level than in the donor atom. The result of the transfer is that the electron-accepting atom becomes a negatively charged ion (an anion), while the other becomes a positive ion (a cation), resulting in an electrostatic attraction between them. Ionic bonds usually occur between metallic atoms (forming cations) and nonmetals (forming anions). The cation and anion bond to form a metal salt, a well known example being sodium chloride, (Na+Cl-).
1.2 States of matter
Solids, liquids and gases
The way atoms interact with one another determines the properties of matter. Interatomic and intermolecular bonds both determine the bulk properties of a compound, including whether it exists as a solid, a liquid or a gas at a given temperature. Solids have rigid bonds between their molecules; liquids have looser bonds; gases have minimal bonds. Table 1.1 summarizes the microscopic properties of solids, liquids and gases. A fourth state of matter: plasma, discussed at the end of this section, can also exist in certain extreme conditions. Figure 1.2 shows the different states of matter and how matter can change from one state to another.
Heat capacity
For an object to increase in temperature, energy in the form of heat must be added and this is covered in more detail in Chapter 4. The specific heat capacity (c) of a substance determines the energy needed to raise 1 kg of the substance by a temperature of 1°C:
Q = c · m · ΔT 1.1
where Q is the energy required
c is the specific heat capacity
m is the mass
ΔT is the temperature change
Water has a specific heat capacity of 4.18 J g-1 °C-1, in other words 4.18 joules are needed to raise 1 g of water by 1°C. Liquid water has a constant heat capacity, regardless of its temperature. The heat capacity of ice is different to that of liquid water, however, which is in turn different to the value for steam.
Latent heat
Suppose you are heating a substance in an oven, say a solid block of wax initially at room temperature. The temperature of the wax will rise steadily as the block absorbs thermal energy (Figure 1.3). When the melting point is reached (around 59°C for paraffin wax), the temperature stops rising even though the wax continues to absorb thermal energy. This energy is used to break the intermolecular bonds, causing the wax to liquefy. Once all the wax has melted, the temperature of the liquid wax will rise once more. If the oven is hot enough the wax will become a vapour and again the temperature will remain constant during the change of phase. For the same reason, water boiling in a kettle remains at 100°C until the water boils away, no matter how high the heat is turned up.
Cooling a hypothermic patient is much more effective if melting ice is used, compared to the same mass of freezing water, even if both are at 0°C. With ice, heat is absorbed from the patient to melt the ice, even though the temperature of the ice does not rise until it has thawed.
For matter to change state, energy must either be added or removed and this energy is referred to as latent heat. When a solid becomes a liquid or when a liquid becomes a gas, energy must be added to break bonds; similarly when bonds are formed, energy is liberated. Latent heat is quantified by the energy required to change the state of one kilogram of matter. The latent heat of vaporization is the energy required to boil one kilogram of liquid. This is the same amount of energy liberated when a kilogram of gas condenses to a liquid (so is also referred to as the latent heat of condensation). Similarly, the latent heat of fusion is the amount of energy liberated when one kilogram of liquid freezes.
Q = m · L 1.2
where Q is the energy required
m is the mass
L is the latent heat
Symbols and units. Latent heat usually has the symbol L along with a subscript: Lf for latent heat of fusion and Lv for latent heat of vaporization. Caution is needed because different terms are used such as latent heat of melting. Latent heat is measured in J·kg-1 or, more commonly, kJ·kg-1 = 1×103 J·kg-1. Specific heat capacity takes the symbol, c, and has the units of J·kg-1·°C-1.
The fourth state of matter: plasma
Plasma is a gas-like mixture of equal numbers of positive and negative ions making it electrically neutral. It is created at very high temperatures when molecules are ripped apart and electrons are stripped from their atoms; the resulting high-energy ions form plasma. Plasma is not encountered under normal conditions, though it should be noted that the sun is composed mainly of matter in the plasma phase. With the aid of electricity, plasma can be used to generate light such as in the fluorescent strip light or a plasma television; a plasma screen consists of thousands of pockets of gases each located between two tiny electrodes.
1.3 Phase diagrams
Temperature has an effect on whether an object is a solid, liquid or gas. Water turns from a liquid to a gas as it boils at 100°C and from liquid to solid as it freezes at 0°C. These temperatures only apply to substances at one particular pressure: atmospheric pressure at sea level. A pan of water on a camping stove near the summit of Mount Everest will boil at less than 80°C as a result of the lower atmospheric pressure at altitude. A pressure cooker maintains higher-than-atmospheric pressure within the cooker, typically allowing water to boil at around 125°C, significantly reducing cooking times.
If the temperature and pressure are known, then a phase diagram can be used to predict what state a substance will be in. Figure 1.4 shows the phase diagram for water, but all substances have a different phase diagram. The point where all three states, solid, liquid and gas, intersect is referred to as the triple point. For water, the triple point is at a temperature of 0.001°C and a pressure of only 0.006 atmospheres. For carbon dioxide, however, the triple point temperature is lower at -56.4°C while the pressure is five times higher than atmospheric pressure.
Above a certain temperature and pressure, known as the critical point, a substance can exist only as a gas, no matter how high the pressure. For water the critical point is 374°C at a pressure of 218 bar. Table 1.2 shows some example elements and compounds with their critical temperatures and pressures.
For water the solid–liquid equilibrium line (the melting point line) slopes backwards rather than forwards. The solid phase, ice, is less dense than the liquid phase, which explains why ice floats in water. It has been postulated that life could not have evolved in the sea if water did not have this highly unusual property.
Is steam a gas?
When water boils it changes state from a liquid to a gas or vapour. 'White steam' consists of tiny droplets of condensed vapour, so is in fact a liquid! True steam is a vapour, invisible to the human eye and is present close to the spout of the kettle.
What is the difference between a gas and a vapour?
The language surrounding gases is confusing, as scientific terms have been mixed in with day-today bywords. A vapour is a type of gas, and is any substance in the gas phase at a temperature lower than its critical temperature. This means that the vapour can be condensed to a liquid or to a solid by increasing the pressure without reducing the temperature. Carbon dioxide, for example, has a critical temperature of 31.03°C, so may be described as a vapour below this temperature.
CHAPTER 2Simple mechanics
Having read this chapter you will be able to:
• Distinguish between force, velocity, speed and acceleration.
• Be familiar with Newton's laws of motion.
• Understand the vector nature of velocity.
• Distinguish between Newtonian and non-Newtonian viscosity.
• List the properties that affect the viscosity of blood.
• Define surface tension and wall tension.
• Apply Laplace's law to alveoli.
• Explain the role played by surfactant.
• Appreciate the critical point in vessel collapse.
2.1 Force, velocity and acceleration
The English philosopher and mathematician Isaac Newton is famous for observing an apple falling to the ground and concluding that all objects with mass are attracted to all others, explaining, among other things, how the moon orbits the earth. This work is embodied in the more general three Laws of Motion, which describe the relationship between velocity, force and acceleration.
Newton's first law: constant velocity
Newton's first law outlines that any change in motion is the result of the application of a force. A ball bearing rolling across a glass surface has almost no friction acting on it so it rolls with almost constant velocity. The Voyager 1 spacecraft left the outer solar system at a speed of 63 000 km·h-1 and will continue to move at this speed forever unless it comes close to a star or other object.
Physics makes a distinction between speed and velocity though they both refer to the rate of travel. Speed is scalar and does not have a direction linked to it whereas velocity is a vector and does include direction. Velocity can be thought of as speed with direction. So if a car drives around a roundabout at a constant 20 km·h-1 then, despite this constant speed, the velocity has changed because the direction of travel has altered. Newton's laws apply to liquids and gases f lowing through a tube; every time a change in direction is required, such as at the junction in a T-piece, forces are needed to produce this change. Fluid f low is examined in more detail in Chapter 8.
Newton's second law: force and acceleration
Newton's second law states that acceleration of a body is proportional to the force applied, and inversely proportional to the mass of the object. Massive objects require more force to generate the same acceleration than small objects. This is why you have to push a loaded shopping trolley much harder than an empty one.
Acceleration is a change in velocity, either speed or direction (or both speed and direction). A tennis ball attached to a string and whirled around your head is accelerating, because the force in the string acts inwards on the ball, causing a change in velocity, despite its speed remaining constant. The earth orbits the sun in exactly the same way, except that gravity replaces the string. Similarly, every twist and turn in the breathing circuit causes the f low of air to change direction, so acceleration of the gas takes place.
(Continues...)
Excerpted from Physics in Anaesthesia by Ben Middleton, Justin Phillips, Rik Thomas, Simon Stacey. Copyright © 2012 Scion Publishing Ltd. Excerpted by permission of Scion Publishing Limited.
All rights reserved. No part of this excerpt may be reproduced or reprinted without permission in writing from the publisher.
Excerpts are provided by Dial-A-Book Inc. solely for the personal use of visitors to this web site.
Table of Contents
1 Atoms and matter
2 Simple mechanics
3 Energy and power
4 Temperature and heat
5 Waves
6 Pressure measurement
7 Humidity
8 Measurement of gas flow
9 The gas laws
10 Diffusion, osmosis and solubility
11 Measuring gas and vapour concentrations
12 Vaporizers
13 Medical gas supplies
14 Breathing systems and ventilation
15 Optics
16 Blood flow measurement
17 Equipment management
18 Basics of electricity
19 Electromagnetism and alternating current
20 Electrical shocks and safety
21 Electrocardiography, pacing and defibrillation
22 Processing, storage and display
23 Ultrasound
24 Lasers
25 Magnetic resonance imaging
26 Nuclear physics and radiation
27 Basic mathematical concepts
28 Physical quantities and SI units
29 Statistics